Sulfur trioxide
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| Sulfur trioxide | |
|---|---|
| Other names | Sulfuric anhydride Sulfan® Sulphur trioxide Sulfur trioxide |
| Identifiers | |
| CAS number | |
| Properties | |
| Molar mass | 80.06 g mol−1 |
| Density | 1.92 g cm−3 |
| Melting point |
16.9 °C, 62.4 °F |
| Boiling point |
45 °C, 113 °F |
| Solubility in other solvents | Hydrolysis |
| Thermochemistry | |
| Std enthalpy of formation ΔfH | −397.77 kJ/mol |
| Standard molar entropy S | 256.77 J.K−1.mol−1 |
| Hazards | |
| EU classification | Corrosive (C) |
| R-phrases | R14, R35, R37 |
| S-phrases | S1/2, S26, S30, S45 |
| Related Compounds | |
| Related compounds | SO2 H2SO4 SO2Cl2 |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references | |
Sulfur trioxide (also spelled sulphur trioxide) is the chemical compound with the formula SO3. In the gaseous form, this species is an important pollutant, being the primary agent in acid rain. It is prepared on massive scale as a precursor to sulfuric acid.
Contents |
Structure and bonding
Gaseous SO3 is a trigonal planar molecule of D3h symmetry, as predicted by VSEPR theory.
In terms of electron-counting formalisms, the sulfur atom has an oxidation state of +6, a formal charge of 0, and is surrounded by 6 electron pairs. From the perspective of molecular orbital theory, most of these electron pairs are non-bonding in character, as is typical for hypervalent molecules.
Chemical reactions
SO3 is the anhydride of H2SO4. Thus, the following reaction occurs:
The reaction occurs both rapidly and exothermically. At or above ~340 °C, sulfuric acid, sulfur trioxide, and water coexist in significant equilibrium concentrations.
Sulfur trioxide also reacts with sulfur dichloride to yield the useful reagent, thionyl chloride.
- SO3 + SCl2 → SOCl2 + SO2
SO3 is a strong Lewis acid readily forming crystalline complexes with pyridine , dioxane and trimethylamine which can be used as sulfonating agents[1].
Preparation
Sulfur trioxide can be prepared in the laboratory by the two-stage pyrolysis of sodium bisulfate:
- 1) dehydration
- 2NaHSO4 → Na2S2O7 + H2O @ 315°C
- 2) cracking
- Na2S2O7 → Na2SO4 + SO3 @ 460°C
This method will work for other metal bisulfates, the controlling factor being the stability of the intermediate pyrosulfate salt.
Industrially SO3 is made by the contact process. Sulfur dioxide, generally made by the burning of sulfur or iron pyrite (a sulfide ore of iron), is first purified by electrostatic precipitation. The purified SO2 is then oxidised by atmospheric oxygen at between 400 and 600 °C over a catalyst consisting of vanadium pentoxide V2O5 activated with potassium oxide K2O on kieselguhr or silica support. Platinum also works very well but is too expensive and is poisoned (rendered ineffective) much more easily by impurities.
The majority of sulphur trioxide made in this way is converted into sulfuric acid not by the direct addition of water, with which it forms a fine mist, but by absorption in concentrated sulfuric acid and dilution with water of the produced oleum.
Structure of solid SO3
The nature of solid SO3 is a surprisingly complex area because of structural changes caused by traces of water.[1] Upon condensation of the gas, absolutely pure SO3 condenses into a trimer, which is often called γ-SO3. This molecular form is a colorless solid with a melting point of 16.8 °C. It adopts a cyclic structure described as [S(=O)2(μ-O)]3[1].
If SO3 is condensed above 27 °C, then α-"SO3" forms, which has a melting point of 62.3°C. α-SO3 is fibrous in appearance, like asbestos (with which it has no chemical relationship). Structurally, it is the polymer [S(=O)2(μ-O)]n. Each end of the polymer is terminated with OH groups (hence α-"SO3" is not really a form of SO3). 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[1].
Relative vapor pressures of solid SO3 are alpha< beta< gamma at identical temperatures, indicative of their relative molecular weights. Liquid sulfur trioxide has vapor pressure consistent with the gamma form. Thus heating a crystal of α-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."[1]
SO3 is aggressively hygroscopic. In fact, the heat of hydration is sufficient that mixtures of SO3 and wood or cotton can ignite. In such cases, SO3 dehydrates these carbohydrates.[1]
Sources
See also
References
ar:ثلاثي أكسيد الكبريت cs:Oxid sírový de:Schwefeltrioxidfr:Trioxyde de soufre it:Triossido di zolfo lt:Sieros trioksidas nl:Zwaveltrioxide ja:三酸化硫黄sk:Oxid sírový sv:Svaveltrioxid vi:SO3 uk:Триоксид сірки
Acknowledgement and Attribution Regarding Sources of Content
Some of the initial content on this page may be incorporated in part from copyleft sources in the public domain including wikis such as Wikipedia and AskDrWiki. Drug information for patients came from the The National Library of Medicine. Infectious disease information may have come from the Centers for Disease Control (CDC). Differential Diagnoses are drawn from clinicians as well as an amalgamation of 3 sources: 1.The Disease Database; 2. Kahan, Scott, Smith, Ellen G. In A Page: Signs and Symptoms. Malden, Massachusetts: Blackwell Publishing, 2004:3; 3. Sailer, Christian, Wasner, Susanne. Differential Diagnosis Pocket. Hermosa Beach, CA: Borm Bruckmeir Publishing LLC, 2002:7 .
