Sulfur dioxide

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Sulfur dioxide
Image:Sulfur-dioxide-2D.png Image:Sulfur-dioxide-3D-vdW.png
General
Systematic name	sulfur dioxide
Other names	Sulfur(IV) oxide Sulfurous anhydride, 'sulphurous anhydride
Molecular formula	S O₂
Molar mass	64.054 g mol⁻¹
Appearance	colourless gas
CAS number	[7446-09-5]
EINECS number	231-195-2
Properties
Density and phase	2.551 g/L, gas
Solubility in water	9.4 g/100 mL (25 °C)
Melting point	−72.4 °C (200.75 K)
Boiling point	−10 °C (263 K)
Critical Point	157.2°C at 7.87 MPa
Acidity (pK_a)	1.81
Structure
Molecular shape	bent
Dipole moment	1.63 D
Thermodynamic data
Standard enthalpy of formation Δ_fH°_gas	−296.84 kJ mol⁻¹
Standard molar entropy S°_gas	248.21 J K⁻¹ mol⁻¹
Safety data
EU classification	Toxic
R-phrases	R23, R34
S-phrases	S1/2, S9, S26 S36/37/39, S45
NFPA 704	Image:NFPA 704.svg 0 3 0
PEL-TWA (OSHA)	5 ppm (13 mg m⁻³)
IDLH (NIOSH)	100 ppm
Flash point	non-flammable
RTECS number	WS4550000
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
Related compounds
Other cations	Selenium dioxide Tellurium dioxide
Related compounds	Sulfur trioxide Sulfuric acid
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references

Sulfur dioxide (also sulphur dioxide) is the chemical compound with the formula SO₂. This important gas is the main product from the combustion of sulfur compounds and is of significant environmental concern. SO₂ is often described as the "smell of burning sulfur."

SO₂ is produced by volcanoes and in various industrial processes. Since coal and petroleum contain various amounts of sulfur compounds, their combustion generates sulfur dioxide. Further oxidation of SO₂, usually in the presence of a catalyst such as NO₂, forms H₂SO₄, and thus acid rain.^{[citation needed]}

1 Preparation
2 Structure and bonding
3 Uses
4 Emissions
5 See also
6 External links
7 Appendix: temperature dependence of aqueous solubility

[edit] Preparation

Sulfur dioxide can be prepared by burning sulfur:

S₈(s) + 8O₂(g) → 8SO₂(g)

The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly.

2H₂S(g) + 3O₂(g) → 2H₂O(g) + 2SO₂(g)

The roasting of sulfide ores such as iron pyrites and sphalerite (zinc blende) also gives SO₂:

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

2ZnS(s) + 3O₂(g) → 2ZnO(s) + 2SO₂(g)

When anhydrous CaSO₄ is heated with coke and sand in the manufacture of cement, CaSiO₃, sulfur dioxide is a by-product.

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

Action of hot concentrated sulphuric acid on copper tunings will produce sulphur dioxide.

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

[edit] Structure and bonding

SO₂ is a bent molecule with C_2v symmetry point group.

In terms of electron-counting formalisms, the sulfur atom has an oxidation state of +4, a formal charge of 0, and is surrounded by 5 electron pairs. From the perspective of molecular orbital theory, most of these electron pairs are non-bonding in character, as is typical for hypervalent molecules.

[edit] Uses

Sulfur dioxide is sometimes used as a preservative in alcoholic drinks, or dried apricots and other dried fruits due to its antimicrobial properties. The preservative is used to maintain the appearance of the fruit rather than prevent rotting. This can give fruit a distinctive chemical taste.

Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances that can be reduced by it; thus making it 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. This might explain why older newspapers turn yellow, because paper used for newspaper is naturally yellow.

Sulfur dioxide is also used to make 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. This is called the contact process.

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.

Sulfur Dioxide blocks nerve signals from the Pulmonary Stretch Receptors (PSR's) and abolishes the Hering-Breuer Inflation Reflex.

Prior to the development of Freons, sulfur dioxide was used as a refrigerant in home refrigerators.

H₂S O₃ is also called "hydrogen sulfite" or sulfurous acid.

[edit] Emissions

According to the US EPA (as presented by the 2002 World Almanac or in chart form [1]), the following amount of sulfur dioxide was released in the U.S. per year, measured in thousands of short tons:

*1999	18,867
*1998	19,491
*1997	19,363
*1996	18,859
*1990	23,678
*1980	25,905
*1970	31,161

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 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 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.

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.[2]

Al-Mishraq, an Iraqi sulfur plant, was the site of a 2004 disaster resulting in the release of massive amounts of sulfur dioxide into the atmosphere.

[edit] See also

[edit] External links

[edit] Appendix: temperature dependence of aqueous solubility

22 g/100ml (0 °C) 15 g/100ml (10 °C) 11 g/100ml (20 °C) 9.4 g/100 ml (25 °C) 8 g/100ml (30 °C) 6.5 g/100ml (40 °C) 5 g/100ml (50 °C) 4 g/100ml (60 °C) 3.5 g/100ml (70 °C) 3.4 g/100ml (80 °C) 3.5 g/100ml (90 °C) 3.7 g/100ml (100 °C)