Sulfur dioxide (IUPAC-recommended spelling) or sulphur dioxide (traditional Commonwealth English) is the chemical compound with the formula SO2. It is a colorless gas with a pungent smell that is responsible for the odor of burnt matches. It is released naturally by volcanic activity and is produced anthropogenically as a by-product of metals refining and the burning of fossil fuels for energy. Sulfur dioxide is somewhat toxic to humans, although only when inhaled in relatively large quantities for a period of several minutes or more. It was known to medieval alchemists as "volatile spirit of sulfur".
Structure and bonding 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.
Occurrence
Sulfur dioxide is found on Earth and exists in very small concentrations in the atmosphere at about 15 ppb. 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 sulfurous acid (SO2 + H2O ⇌ HSO−3 + H+), and is a key component of the planet's global atmospheric sulfur cycle. 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 and trace amounts are thought to also exist in the atmosphere of Jupiter. The James Webb Space Telescope has observed the presence of sulfur dioxide on the exoplanet WASP-39b, where it is formed through photochemistry in the planet's atmosphere. 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 (see contact process, but other processes predated that at least since 16th century). In the United States in 1979, 23.6 million metric tons (26 million U.S. short tons) of sulfur dioxide were used in this way, compared with 150,000 metric tons (165,347 U.S. 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:
S8 + 8 O2 → 8 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 1,000–1,600 °C (1,830–2,910 °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 SO2 + 2 H2O The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases SO2:
4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2 2 ZnS + 3 O2 → 2 ZnO + 2 SO2 HgS + O2 → Hg + SO2 4 FeS + 7 O2 → 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 SO2.
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 Tin also reacts with concentrated sulfuric acid but it produces tin(II) sulfate which can later be pyrolyzed at 360 °C into tin dioxide and dry sulfur dioxide.
Sn + H2SO4 → SnSO4 + H2 SnSO4 → SnO2 + SO2
From sulfites The reverse reaction occurs upon acidification:
H+ + HSO−3 → SO2 + H2O
Reactions Sulfites result by the action of aqueous base on sulfur dioxide:
SO2 + 2 NaOH → Na2SO3 + H2O Sulfur dioxide is a mild but useful reducing agent. It is oxidized by halogens to give the sulfuryl halides, such as sulfuryl chloride:
SO2 + Cl2 → SO2Cl2 Sulfur dioxide is the oxidising agent in the Claus process, which is conducted on a large scale in oil refineries. Here, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:
SO2 + 2 H2S → 3 S + 2 H2O The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid.
2 SO2 + 2 H2O + O2 → 2 H2SO4 Sulfur dioxide dissolves in water to give "sulfurous acid", which cannot be isolated and is instead an acidic solution of bisulfite, and possibly sulfite, ions.
SO2 + H2O ⇌ HSO−3 + H+ Ka = 1.54×10−2; pKa = 1.81
Laboratory reactions Sulfur dioxide is one of the few common acidic yet reducing gases. It turns moist litmus pink (being acidic), then white (due to its bleaching effect). It may be identified by bubbling it through a dichromate solution, turning the solution from orange to green (Cr3+ (aq)). It can also reduce ferric ions to ferrous. Sulfur dioxide can react with certain 1,3-dienes in a cheletropic reaction to form cyclic sulfones. This reaction is exploited on an industrial scale for the synthesis of sulfolane, which is an important solvent in the petrochemical industry.
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