Sulfur dioxide

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)

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Sulfur dioxide
Other names	Sulfur(IV) oxide Sulfurous anhydride
Molecular formula	S O₂
Molar mass	64.06 g/mol
Appearance	colourless gas
CAS number	^*
EINECS number	231-195-2
Properties
Density and phase	2.551 g/L, gas
Solubility in water	22 g/100ml (0 °C)
Melting point	−72.4 °C (200.75 K)
Boiling point	−10 °C (263 K)
Acidity (pK_a)	1.81
Structure
Molecular shape	bent
Dipole moment	1.63 D
Thermodynamic data
Standard enthalpy of	−296.84 kJ/mol
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
PEL-TWA (OSHA)	5 ppm (13 mg/m³)
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) Chemical infobox

Sulfur dioxide (also sulphur dioxide, sulfurous anhydride or sulphurous anhydride) has the chemical formula S O₂. The gas is irritating to the lungs and is frequently described as smelling of burning sulfur.

It is produced by volcanoes and in various industrial processes. In particular, poor-quality coal and petroleum contain sulfur compounds, and generate sulfur dioxide when burned: the gas reacts with water and atmospheric oxygen to form sulfuric acid (H₂S O₄) and thus acid rain.

Preparation

Sulfur dioxide is often prepared by burning sulfur in air:

S_(s) + O_{2 (g)} → S O_{2 (g)}

Hydrogen sulfide from crude oil may also be burned.

2H₂S_(g) + 3O_{2 (g)} → 2H₂O_(g) + 2S O_{2 (g)}

Sulfide ores such as iron pyrites and sphalerite (zinc blende) may also be used:

4Fe S_{2 (s)} + 11O_{2 (g)} → 2Fe₂O_{3 (s)} + 8SO_2 (g)

2Zn S _(s) + 3O_{2 (g)} → 2Zn O _(s) + 2S O_{2 (g)}

When anhydrous Ca S O₄, is heated with coke and sand in the manufacture of cement, Ca Si O₃, sulfur dioxide is a by-product.

2Ca S O_{4 (s)} + 2Si O_{2 (s)} + C _(s) → 2Ca Si O_{3 (s)} + 2S O_{2 (g)} + C O_{2 (g)}

Uses

Sulfur dioxide is sometimes used as a preservative in alcoholic drinks, or dried apricots and other dried fruits. 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.

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

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

Emissions

According to the EPA (as presented by the 2002 World Almanac or in chart form ^*), the following amount of thousands of short tons of Sulfur dioxide were released in the U.S. per year:

|*1999

|*1998

|*1997

|*1996

|*1990

|*1980

|*1970

18,867

19,491

19,363

18,859

23,678

25,905

31,161

Due largely to EPA’s Acid Rain Program there has been a 33 percent decrease in emissions between 1983 and 2002, partly due to flue gas desulfurization.

In power plants using sulfur containing coal or oil the sulfur dioxide can be chemically bound by flue gas desulfurization . Calcium oxide reacts with sulfur dioxide to form calcium sulfite. Most of the gypsum sold in Europe comes from this source.

Ca O + S O₂ → Ca S O₃

External links

Gourt Home::Gourt :: Sulfur dioxide

Preparation

Uses

Emissions

See also

External links