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SULFUR DIOXIDE
| Sulfur dioxide |
  |
| General |
| Systematic name |
sulfur dioxide |
| Other names |
Sulfur(IV) oxide
Sulfurous anhydride |
| Molecular formula |
SO2 |
| Molar mass |
64.06 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 |
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)
|
| Melting point |
−72.4 °C (200.75 K) |
| Boiling point |
−10 °C (263 K) |
| Critical Point |
157.2°C at 7.87 MPa |
| Acidity (pKa) |
1.81 |
| Structure |
| Molecular shape |
bent |
| Dipole moment |
1.63 D |
| Thermodynamic data |
Standard enthalpy
of formation ΔfH°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 |
|
| 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, εr, 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, sulfurous anhydride or sulphurous anhydride) has the chemical formula SO2. 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 (H2SO4) and thus acid rain.
Preparation
Sulfur dioxide is often prepared by burning sulfur in air:
- S(s) + O2(g) → SO2(g)
Hydrogen sulfide from crude oil may also be burned.
- 2H2S(g) + 3O2(g) → 2H2O(g) + 2SO2(g)
Sulfide ores such as iron pyrites and sphalerite (zinc blende) may also be used:
- 4FeS2(s) + 11O2(g) → 2Fe2O3(s) + 8SO2(g)
- 2ZnS(s) + 3O2(g) → 2ZnO(s) + 2SO2
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)
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 Freons, sulfur dioxide was used as a refrigerant in home refrigerators.
H2SO3 is also called "hydrogen sulfite" or sulfurous acid.
Emissions
According to the US EPA (as presented by the 2002 World Almanac or in chart form [1]), the following amount of thousands of short tons of Sulfur dioxide were released in the U.S. per year:
| *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 improved 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.
See also
External links
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