Oxides of sulfur

Sulfur dioxide is manufactured on a large scale by burning sulfur (the most important process) or H2S, by roasting sulfide ores (e.g. equation 16.85), or reducing CaS04 (equation 16.86). Desulfurization processes to limit SO2 emissions (see Box 12.2) and reduce acid rain (see Box 16.5) are now in use. In the laboratory, SO2 may be prepared by, for example, reaction 16.87, and it is commercially available in cylinders. Selected physical properties of SO2 are listed in Table 16.7. 

The boiling point of SO2 is 263 K, but it can be safely handled in a sealed tube at room temperature under its own vapour pressure. It is a good solvent with a wide range of uses (see Section 9.8). Sulfur dioxide has a molecular structure (16.44). 

Physical appearance and general characteristics Melting point / K Boiling point/K Avap7/°(bp) / kJ moP  

By the time acid rain falls to the Earth s surface, the pollutants may have travelled long distances from their industrial sources. For example, prevailing winds in Europe may carry SO2 from the UK, France and Germany to Scandinavia. 

International legislation to reduce acidic gas emissions has been in operation since the 1980s. The graph on the next page illustrates the trend in emissions of SO2 and NO2 in Europe from 1880 and projected to 2020, and this 

Sulfur dioxide reacts with O2 (see below), F2 and CI2 (eq. 16.88). It also reacts with the heavier alkah metal fluorides to give metal fluorosulfites (eq. 16.89), and with CSN3 to give the Cs salts of [S02N3] (Fig. 16.16a) and [(S02)2N3] (Fig. 16.16b). The latter is formed when CsNa dissolves in liquid SO2 at 209 K. On raising the temperature to 243 K, [(S02)2N3] loses one equivalent of SO2 to yield [SO2N3]-. 

In aqueous solution, SO2 is converted to only a small extent to sulfurous acid. Aqueous solutions of H2SO3 contain significant amounts of dissolved SO2 (see eqs. 7.18-7.20). Sulfur dioxide is a weak reducing agent in acidic solution, and a slightly stronger one in basic media (eqs. 16.90 and 16.91). 

More than ten oxides of sulfur are known, and the most important industrially are SO2 and SO3. The six homocyclic polysulfur monoxides S 0 (5 n 10) are prepared by oxidizing the appropriate cyclo-S with trifluoroperoxoacetic 

Coordination modes of SO2 to metal atoms (shaded circles). 

When sulfide ores (such as pyrite), sulfur-rich organic compounds, and fossil fuels (such as coal) are burned in air, the sulfur therein is mostly converted to SO2. 

Sulfur dioxide is a colorless, toxic gas with a choking odor. Gaseous SO2 neither burns nor supports combustion. It is readily soluble in water (3927 cm3 SO2 in 100 g H2O at 20°C) and forms sulfurous acid.  

Sulfur dioxide is a component of air pollutants, and is capable of causing severe damage to human and other animal lungs, particularly in the sulfate form. It is also an important precursor to acid rain. The SO2 molecule survives for a few days in the atmosphere before it is oxidized to SO3. The direct reaction 

All these compounds have (distorted) tetrahedral molecules, those of formula O2SX2 having C2v symmetry and the others Cj. Dimensions are in Table 15.15 the remarkably short O-S and S-F distances in O2SF2 should be noted (cf. above). Indeed, the implied strength of bonding in this molecule is reflected by the fact that it can be made by reacting the normally extremely inert compound SFg (p. 687) with the fluoro-acceptor SO3  

Gmelin Handbuch der Anorganischen Chemie, 8th edn., Schwefel Oxide, Erganzungsband 3, 1980, 344 pp. 

S7O2 dark orange crystals, dccomp room temp 

The fugitive species SO was first identified by its ultraviolet spectrum in 1929 but it is thermodynamically unstable and decomposes completely in the gas phase in less than I s. It is formed by reduction of SOn with sulfur vapour in a glow discharge and its spectroscopic properties 

4-centre-2 electron (4c-2e) as in [Fe3(CO)gS(/r3-SO)],t Mi 2c-2e as in (IrCl(SO)(PR,)2l, and also 3c-4e and 3c-2e in several dinu-clear transition-metal complexes.A novel and unprecedented route to this last class of fi-SO complexes involves the direct oxidative addition of OSCF to the Ni complex [Ni(cod)21 in the presence of dppm (cod = cycloocta-1,5-diene, dppm = Ph2PCH2PPh2) to form the purple crystalline dinickel A-frame complex, [Ni2(/r-SO)(dppm)2Cl2. X-ray analysis reveals two slightly differing geometries, with SO 

144 and 145.9 pm, respectively (both shorter than in the free molecule, 148.1 pm), and with the SO ligand being tilted with respect to the Ni - Ni 

The most important oxides of sulfur are SO2 and SO3, but there are also a number of unstable oxides. Among these 

The effects of acid rain can be devastating. The pH of lakes and streams is lowered, although the composition of the bedrock is significant, and in some cases provides a natural buffering effect. A second effect is that acid rain penetrating the bedrock can react with aluminosilicate minerals, or can 

The sulfur-oxygen bond is an important feature of sulfur chemistry, just as the silicon-oxygen bond is of silicon chemistry. The best-known simple oxides are the dioxide, SO2, and the trioxide, SO3 (although highly reactive lower oxides SO and S2O have been described). The bonding in SO2 and SO3 is usually described in terms of S=0 double bonds, VSEPR theory predicting the shapes shown in Structures 12.14-12.17. 

with its characteristic choking smell, is well known as the product of combustion of sulfur or sulfides in air or oxygen. SO2 is readily soluble in water, dissolving to form sulfurous acid (an acid that has never been isolated as a pure compound) and the sulfite ion  

As a consequence, sulfur dioxide in the presence of water is an effective reducing agent, supplying electrons to an oxidizing agent, via the sulfite ion  

This can easily be demonstrated visually by bubbling sulfur dioxide into an aqueous solution of acidified potassium dichromate. There is a rapid colour change from orange to green as the chromium(VI) is reduced to chromium(III)  

The reducing properties of sulfur dioxide are put to good use commercially in bleaching natural fibres (straw, wool and newsprint). Sulfur dioxide has long been used as an antiseptic and antioxidant in the food industry wine casks have been fumigated with SO2 for thousands of years. The home winemaker of today, who adds a sodium metabisulfite tablet to the wash water, is following a very ancient practice The sodium metabisulfite dissolves to release SO2, which sterilizes the equipment. 

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Utility systems as sources of waste. The principal sources of utility waste are associated with hot utilities (including cogeneration systems) and cold utilities. Furnaces, steam boilers, gas turbines, and diesel engines all produce waste from products of combustion. The principal problem here is the emission of carbon dioxide, oxides of sulfur and nitrogen, and particulates (metal oxides, unbumt... 

The electrolytic processes for commercial production of hydrogen peroxide are based on (/) the oxidation of sulfuric acid or sulfates to peroxydisulfuric acid [13445-49-3] (peroxydisulfates) with the formation of hydrogen and (2) the double hydrolysis of the peroxydisulfuric acid (peroxydisulfates) to Caro s acid and then hydrogen peroxide. To avoid electrolysis of water, smooth platinum electrodes are used because of the high oxygen overvoltage. The overall reaction is... 

S-oxidation of sulfur-containing pesticides such as aldicarb, parathion, and malathion can be of importance in the absence of microbial activity (29). The products of chemical vs biological oxidation are generally identical (eq. 8). 

Numerous oxides of sulfur have been reported and those that have been characterized are SO [13827-32-2] S2O [20901 -21 -7] S O (n = 6-10), SO2, SO, and SO4 [12772-98-4]. Among these, SO2 and SO ate of principal importance. Sulfur oxide chemistry has been reviewed (210—212). Sulfur trioxide, SO, is discussed elsewhere (see Sulfuric acid and sulfur trioxide). 

Oxidation of sulfur dioxide in aqueous solution, as in clouds, can be catalyzed synergistically by iron and manganese (225). Ammonia can be used to scmb sulfur dioxide from gas streams in the presence of air. The product is largely ammonium sulfate formed by oxidation in the absence of any catalyst (226). The oxidation of SO2 catalyzed by nitrogen oxides was important in the eady processes for manufacture of sulfuric acid (qv). Sulfur dioxide reacts with chlorine or bromine forming sulfuryl chloride or bromide [507-16 ]. 

The anaerobic reaction of sulfur dioxide with aqueous ammonia produces a solution of ammonium sulfite [10192-30-0]. This reaction proceeds efficientiy, even with a gas stream containing as Httie as 1 wt % sulfur dioxide. The sulfur dioxide can be regenerated at a high concentration by acidulation or by stream stripping of the ammonium sulfite solution, or the sulfite can be made to precipitate and the ammonia recovered by addition of lime (243). The process can also be modified to produce ammonium sulfate for use as fertili2er (244) (see Fertilizers). In a variant of this process, the use of electron-beam radiation cataly2es the oxidation of sulfur dioxide in the presence of ammonia to form ammonium sulfate (245). 

Solutions of iron chelates can be used to remove hydrogen sulfide and oxides of sulfur and nitrogen in industrial gas scmbbing processes (41,50,51) before flue gases are released to the atmosphere. 

Oxides of sulfur-containing azoles comprise another class of non-aromatic azoles. 

The problems with the combustion reaction occur because the process also produces many other products, most of which are termed air pollutants. These can be carbon monoxide, carbon dioxide, oxides of sulfur, oxides of nitrogen, smoke, fly ash, metals, metal oxides, metal salts, aldehydes, ketones, acids, polynuclear hydrocarbons, and many others. Only in the past few decades have combustion engineers become concerned about... 

The most widely used pulping process is the kraft process, as shown in Fig. 6-11, which results in recovery and regeneration of the chemicals. This occurs in the recovery furnace, which operates with both oxidizing and reducing zones. Emissions from such recovery furnaces include particulate matter, very odorous reduced sulfur compounds, and oxides of sulfur. If extensive and expensive control is not exercised over the kraft pulp process, the odors and aerosol emissions will affect a wide area. Odor complaints have been reported over 100 km away from these plants. A properly controlled and operated kraft plant will handle huge amounts of material and produce millions of kilograms of finished products per day, with little or no complaint regarding odor or particulate emissions. 

The principal components of atmospheric chemical processes are hydrocarbons, oxides of nitrogen, oxides of sulfur, oxygenated hydrocarbons, ozone, and free radical intermediates. Solar radiation plays a crucial role in the generation of free radicals, whereas water vapor and temperature can influence particular chemical pathways. Table 12-4 lists a few of the components of each of these classes. Although more extensive tabulations may be found in "Atmospheric Chemical Compounds" (8), those listed in... 

Oxides of sulfur Sulfur dioxide SO2 Sulfur trioxide SO3 Sulfuric acid H2SO4 Ammonium bisulfate (NH))HS04 Ammonium sulfate (NH4)2S04... 

Oxides of sulfur react with copper oxide to form copper sulfate. Removal with a dry particulate control system follows. 

Why are "oxides of nitrogen" and "oxides of sulfur" usually reported in emission inventory tables rather than the actual oxidation states ... 

Sulfur oxides (SO,) are compounds of sulfur and oxygen molecules. Sulfur dioxide (SO2) is the predominant form found in the lower atmosphere. It is a colorless gas that can be detected by taste and smell in the range of 1, (X)0 to 3,000 uglm. At concentrations of 10,000 uglm , it has a pungent, unpleasant odor. Sulfur dioxide dissolves readily in water present in the atmosphere to form sulfurous acid (H SOj). About 30% of the sulfur dioxide in the atmosphere is converted to sulfate aerosol (acid aerosol), which is removed through wet or dry deposition processes. Sulfur trioxide (SO3), another oxide of sulfur, is either emitted directly into the atmosphere or produced from sulfur dioxide and is readily converted to sulfuric acid (H2SO4). 

Fire Hazards - Flash Point Not flammable Flammable Limits in Air (%) Not flammable Fire Extinguishing AgerUs Not pertinent Fire Extinguishing Agents Not To Be Used Not pertinent Special Hazards of Combustion Products Oxides of sulfur and phosphorous may be formed when exposed to a fire situation Behavior in Fire Data not available Ignition Temperature Not pertinent Electrical Hazard Not pertinent Bunting Rate Not pertinent. 

Used Not pertinent Special Hazards of Combustion Products Irritating oxides of sulfur may be formed Behavior in Fire Not pertinent Ignition Temperature No data Electrical Hazard Not pertinent Burning Rate Not pertinent. 

Vanadium pentoxide, VjOj, is used as a eatalyst in the oxidation of sulfur dioxide. The meehanism involves oxidation-reduetion of V2O5 that exists on the support at operating eonditions in the molten state. The meehanism of reaetion is ... 

The gases normally monirored include oxides of sulfur and nitrogen (SO, NO ),... 

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