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| Names | |
|---|---|
| Preferred IUPAC name Sulfur trioxide | |
| Systematic IUPAC name Sulfonylideneoxidane | |
| Other names Sulfuric anhydride, Sulfur(VI) oxide | |
| Identifiers | |
3D model (JSmol) |
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| ChEBI | |
| ChemSpider |
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| ECHA InfoCard | 100.028.361 |
| EC Number |
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| 1448 | |
| RTECS number |
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| UNII | |
| UN number | UN 1829 |
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| Properties | |
| SO3 | |
| Molar mass | 80.066 g/mol |
| Appearance | Colorless to white crystalline solid which will fume in air.[2] Colorless liquid and gas.[3] |
| Odor | Varies. Vapor is pungent; like sulfur dioxide.[4] Mist is odorless.[3] |
| Density | 1.92 g/cm3, liquid |
| Melting point | 16.9 °C (62.4 °F; 290.0 K) |
| Boiling point | 45 °C (113 °F; 318 K) |
| Reacts to givesulfuric acid | |
| Thermochemistry | |
Std molar entropy(S⦵298) | 256.77 JK−1mol−1 |
Std enthalpy of formation(ΔfH⦵298) | −395.7 kJ/mol |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards | Highly corrosive, extremely strong dehydrating agent |
| GHS labelling: | |
  | |
| Danger | |
| H314,H335 | |
| P261,P280,P305+P351+P338,P310[5] | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LC50 (median concentration) | rat, 4 hr 375 mg/m3[citation needed] |
| Safety data sheet (SDS) | ICSC 1202 |
| Related compounds | |
Othercations | Selenium trioxide Tellurium trioxide Polonium trioxide |
| Sulfur monoxide Sulfur dioxide | |
Related compounds | Sulfuric acid |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Sulfur trioxide (alternative spellingsulphur trioxide) is the chemical compound with the formula SO3. It has been described as "unquestionably the most [economically] importantsulfur oxide".[1] It is prepared on an industrial scale as aprecursor tosulfuric acid.
Sulfur trioxide exists in several forms: gaseous monomer, crystalline trimer, and solid polymer. Sulfur trioxide is a solid at just below room temperature with a relatively narrow liquid range. Gaseous SO3 is the primary precursor toacid rain.[6]
The molecule SO3 istrigonal planar. As predicted byVSEPR theory, its structure belongs to the D3hpoint group. The sulfur atom has anoxidation state of +6 and may be assigned aformal charge value as low as 0 (if all three sulfur-oxygen bonds are assumed to be double bonds) or as high as +2 (if theOctet Rule is assumed).[7] When the formal charge is non-zero, the S-O bonding is assumed to be delocalized. In any case the three S-Obond lengths are equal to one another, at 1.42 Å.[1] The electricaldipole moment of gaseous sulfur trioxide is zero.
Both liquid and gaseous[8] SO3 exists in an equilibrium between the monomer and the cyclic trimer. The nature of solid SO3 is complex and at least 3polymorphs are known, with conversion between them being dependent on traces of water.[9]
Absolutely pure SO3 freezes at 16.8 °C to give theγ-SO3 form, which adopts the cyclic trimer configuration [S(=O)2(μ-O)]3.[10][1]
If SO3 is condensed above 27 °C, thenα-SO3 forms, which has a melting point of 62.3 °C.α-SO3 is fibrous in appearance. Structurally, it is thepolymer [S(=O)2(μ-O)]n. Each end of the polymer is terminated with OH groups.[1]β-SO3, like the alpha form, is fibrous but of different molecular weight, consisting of an hydroxyl-capped polymer, but melts at 32.5 °C. Both the gamma and the beta forms are metastable, eventually converting to the stable alpha form if left standing for sufficient time. This conversion is caused by traces of water.[11]
Relative vapor pressures of solid SO3 are alpha < beta < gamma at identical temperatures, indicative of their relativemolecular weights. Liquid sulfur trioxide has a vapor pressure consistent with the gamma form. Thus heating a crystal ofα-SO3 to its melting point results in a sudden increase in vapor pressure, which can be forceful enough to shatter a glass vessel in which it is heated. This effect is known as the "alpha explosion".[11]
Sulfur trioxide undergoes many reactions.[1]
SO3 is theanhydride of H2SO4. Thus, it is susceptible to hydration:
Gaseous sulfur trioxide fumes profusely even in a relatively dry atmosphere owing to formation of a sulfuric acid mist.SO3 is aggressivelyhygroscopic. The heat of hydration is sufficient that mixtures of SO3 and wood or cotton can ignite. In such cases, SO3 dehydrates thesecarbohydrates.[11]
Akin to the behavior of H2O,hydrogen fluoride adds to givefluorosulfuric acid:
SO3 reacts with dinitrogen pentoxide to give thenitronium salt of pyrosulfate:
Sulfur trioxide is anoxidant. It oxidizessulfur dichloride tothionyl chloride.
SO3 is a strongLewis acid readily forming adducts with Lewis bases.[13] Withpyridine, it gives thesulfur trioxide pyridine complex. Related adducts form fromdioxane andtrimethylamine.
Sulfur trioxide is a potentsulfonating agent, i.e. it adds SO3 groups to substrates. Often the substrates are organic, as inaromatic sulfonation.[14] For activated substrates, Lewis base adducts of sulfur trioxide are effective sulfonating agents.[15]
The direct oxidation of sulfur dioxide to sulfur trioxide in air proceeds very slowly:
Industrially SO3 is made by thecontact process.Sulfur dioxide is produced by the burning ofsulfur oriron pyrite (a sulfide ore of iron). After being purified by electrostatic precipitation, the SO2 is then oxidised by atmosphericoxygen at between 400 and 600 °C over a catalyst. A typical catalyst consists ofvanadium pentoxide (V2O5) activated withpotassium oxide K2O onkieselguhr orsilica support.Platinum also works very well but is too expensive and is poisoned (rendered ineffective) much more easily by impurities.[16]The majority of sulfur trioxide made in this way is converted intosulfuric acid.
Sulfur trioxide can be prepared in the laboratory by the two-stagepyrolysis ofsodium bisulfate.Sodium pyrosulfate is an intermediate product:[17]
The latter occurs at much lower temperatures (45–60 °C) in the presence of catalyticH2SO4.[18] In contrast, KHSO4 undergoes the same reactions at a higher temperature.[17]
Another two step method involving a salt pyrolysis starts with concentrated sulfuric acid and anhydrous tin tetrachloride:
To further reduce water contamination, Oleum and a slight excess of Tin(IV) Chloride should be used. The slight excess of SnCl4 can then be separated by carefully heating the solid Tin(IV) Sulfate under a vacuum to no more than 120 °C. The excess SO3 from the Oleum and the remaining SnCl4 will react during HCl formation and form Tin(IV) Oxide and Sulfuryl Chloride. If an excess of SO3 in the Oleum is present relative to SnCl4 , the Tin(IV) Oxide will absorb it and form more Tin(IV) Sulfate.
The advantage of this method over the sodium bisulfate one is that it can produce the pure trimer of SO3 (since no water is present) while still using safe temperatures for normal borosilicate laboratory glassware. Other dry sulfate salt pyrolysis reactions require higher temperatures which increases the risk of shattering. A disadvantage is that it generates significant quantities of hydrogen chloride gas which needs to be captured as well.
SO3 may also be prepared by dehydratingsulfuric acid withphosphorus pentoxide.[19]
Sulfur trioxide is a reagent insulfonation reactions.Dimethyl sulfate is produced commercially by the reaction ofdimethyl ether with sulfur trioxide:[20]
Sulfate esters are used asdetergents,dyes, andpharmaceuticals. Sulfur trioxide is generatedin situ from sulfuric acid or is used as a solution in the acid.
B2O3 stabilized sulfur trioxide was traded by Baker & Adamson under the tradename "Sulfan" in the 20th century.[21]
Along with being an oxidizing agent, sulfur trioxide is highly corrosive. It reacts violently with water to produce highly corrosive sulfuric acid.
Sulfur trioxide used was pure, colorless liquid SO3 marketed under the trade name Sulfan by Baker and Adamson
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