Difference between revisions of "Sulfur trioxide"
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| OtherNames = Sulfuric anhydride<br/>Sulfan | | OtherNames = Sulfuric anhydride<br/>Sulfan | ||
| Section1 = {{Chembox Identifiers | | Section1 = {{Chembox Identifiers | ||
+ | | InChI = 1/O3S/c1-4(2)3 | ||
+ | | StdInChI = 1S/O3S/c1-4(2)3 | ||
+ | | InChIKey = AKEJUJNQAAGONA-UHFFFAOYAX | ||
+ | | StdInChIKey = AKEJUJNQAAGONA-UHFFFAOYSA-N | ||
| CASNo = 7446-11-9 | | CASNo = 7446-11-9 | ||
| CASNo_Ref = {{cascite}} | | CASNo_Ref = {{cascite}} | ||
+ | | EC-number = 231-197-3 | ||
+ | | ChemSpiderID = 23080 | ||
| UNNumber = 1829 | | UNNumber = 1829 | ||
| RTECS = WT4830000 | | RTECS = WT4830000 | ||
Line 21: | Line 27: | ||
| MeltingPt = 16.9 °C | | MeltingPt = 16.9 °C | ||
| BoilingPt = 45 °C | | BoilingPt = 45 °C | ||
− | | Solubility = | + | | Solubility = reacts violently |
}} | }} | ||
| Section4 = {{Chembox Thermochemistry | | Section4 = {{Chembox Thermochemistry | ||
− | | DeltaHf = | + | | DeltaHf = –397.77 kJ/mol |
− | | Entropy = 256.77 J K<sup> | + | | Entropy = 256.77 J K<sup>–1</sup> mol<sup>–1</sup>}} |
| Section7 = {{Chembox Hazards | | Section7 = {{Chembox Hazards | ||
− | | | + | | Reference = <ref>{{CLP Regulation|index=016-019-00-2|page=400}}</ref> |
− | | | + | | ExternalMSDS = {{ICSC-small|12|02}} |
− | | | + | | EUIndex = 016-019-00-2 <!-- as oleum --> |
− | | | + | | GHSPictograms = {{GHS05|Skin Corr. 1A}}{{GHS07|STOT SE 3}} |
− | | | + | | GHSSignalWord = DANGER |
− | | | + | | HPhrases = {{H-phrases|314|335}} <!-- also EUH014 in the European Union --> |
− | | | ||
− | | | ||
− | |||
| FlashPt = Non-flammable | | FlashPt = Non-flammable | ||
− | |||
− | |||
}} | }} | ||
| Section8 = {{Chembox Related | | Section8 = {{Chembox Related | ||
− | |||
| OtherCations = [[Selenium trioxide]]<br/>[[Tellurium trioxide]] | | OtherCations = [[Selenium trioxide]]<br/>[[Tellurium trioxide]] | ||
| OtherFunctn = [[Sulfur monoxide]]<br/>[[Sulfur dioxide]] | | OtherFunctn = [[Sulfur monoxide]]<br/>[[Sulfur dioxide]] | ||
| Function = [[sulfur]] [[oxide]]s | | Function = [[sulfur]] [[oxide]]s | ||
− | | OtherCpds = [[Sulfuric acid]] | + | | OtherCpds = [[Sulfuric acid]]<br/>[[Disulfuric acid]] |
}} | }} | ||
}} | }} | ||
− | '''[[Sulfur]] trioxide''' (also spelled '''sulphur trioxide''') is the chemical compound with the formula SO<sub>3</sub>. | + | '''[[Sulfur]] trioxide''' (also spelled '''sulphur trioxide''') is the chemical compound with the formula SO<sub>3</sub>. It is prepared industrially on massive scales as a precursor to [[sulfuric acid]]. |
==Structure and bonding== | ==Structure and bonding== | ||
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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 molecule]]s. | 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 molecule]]s. | ||
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==Chemical reactions== | ==Chemical reactions== | ||
SO<sub>3</sub> is the [[anhydride]] of H<sub>2</sub>SO<sub>4</sub>. Thus, the following reaction occurs: | SO<sub>3</sub> is the [[anhydride]] of H<sub>2</sub>SO<sub>4</sub>. Thus, the following reaction occurs: | ||
+ | :SO<sub>3</sub> + H<sub>2</sub>O → H<sub>2</sub>SO<sub>4</sub> (–88 kJ/mol) | ||
− | + | The reaction occurs both rapidly and exothermically, too violently to be used in large-scale manufacturing. Above 340 °C, sulfuric acid, sulfur trioxide, and water coexist in significant equilibrium concentrations. | |
− | |||
− | The reaction occurs both rapidly and exothermically, too violently to be used in large-scale manufacturing. | ||
Sulfur trioxide also reacts with [[sulfur dichloride]] to yield the useful [[reagent]], [[thionyl chloride]]. | Sulfur trioxide also reacts with [[sulfur dichloride]] to yield the useful [[reagent]], [[thionyl chloride]]. | ||
− | |||
:SO<sub>3</sub> + SCl<sub>2</sub> → SOCl<sub>2</sub> + SO<sub>2</sub> | :SO<sub>3</sub> + SCl<sub>2</sub> → SOCl<sub>2</sub> + SO<sub>2</sub> | ||
− | SO<sub>3</sub> is a strong Lewis acid readily forming crystalline complexes with [[sulfur trioxide pyridine complex|pyridine]], [[dioxane]] and [[trimethylamine]] which can be used as sulfonating agents.<ref>{{Cotton&Wilkinson6th}}</ref> | + | SO<sub>3</sub> is a strong Lewis acid readily forming crystalline complexes with [[sulfur trioxide pyridine complex|pyridine]], [[dioxane]] and [[trimethylamine]], which can be used as sulfonating agents.<ref>{{Cotton&Wilkinson6th}}</ref> |
==Preparation== | ==Preparation== | ||
− | Sulfur trioxide can be prepared in the laboratory by the two-stage [[pyrolysis]] of [[sodium bisulfate]]. | + | Sulfur trioxide can be prepared in the laboratory by the two-stage [[pyrolysis]] of [[sodium bisulfate]]. [[Sodium pyrosulfate]] is an intermediate product: |
− | |||
<ol> | <ol> | ||
<li>Dehydration at 315°C: | <li>Dehydration at 315°C: | ||
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This method will work for other metal bisulfates, the controlling factor being the stability of the intermediate pyrosulfate salt. | This method will work for other metal bisulfates, the controlling factor being the stability of the intermediate pyrosulfate salt. | ||
− | Industrially SO<sub>3</sub> 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 SO<sub>2</sub> is then oxidised by atmospheric [[oxygen]] at between 400 and 600 °C over a catalyst consisting of [[Vanadium(V) oxide|vanadium pentoxide]] (V<sub>2</sub>O<sub>5</sub>) activated with [[potassium oxide]] K<sub>2</sub>O on [[Diatomaceous earth|kieselguhr]] or [[Silicon dioxide|silica]] support. [[Platinum]] also works very well but is too expensive and is poisoned (rendered ineffective) much more easily by impurities. | + | Industrially SO<sub>3</sub> 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 SO<sub>2</sub> is then oxidised by atmospheric [[oxygen]] at between 400 and 600 °C over a catalyst consisting of [[Vanadium(V) oxide|vanadium pentoxide]] (V<sub>2</sub>O<sub>5</sub>) activated with [[potassium oxide]] K<sub>2</sub>O on [[Diatomaceous earth|kieselguhr]] or [[Silicon dioxide|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]]. | 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]]. | ||
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==Structure of solid SO<sub>3</sub>== | ==Structure of solid SO<sub>3</sub>== | ||
[[File:Sulfur-trioxide-trimer-from-xtal-1967-3D-balls-B.png|thumb|right|200px|[[Ball-and-stick model]] of the ''γ''-SO<sub>3</sub> molecule]] | [[File:Sulfur-trioxide-trimer-from-xtal-1967-3D-balls-B.png|thumb|right|200px|[[Ball-and-stick model]] of the ''γ''-SO<sub>3</sub> molecule]] | ||
− | The nature of solid SO<sub>3</sub> is a surprisingly complex area because of structural changes caused by traces of water.<ref>{{Holleman&Wiberg}}</ref> Upon condensation of the gas, absolutely pure SO<sub>3</sub> condenses into a trimer, which is often called | + | The nature of solid SO<sub>3</sub> is a surprisingly complex area because of structural changes caused by traces of water.<ref>{{Holleman&Wiberg}}.</ref> Upon condensation of the gas, absolutely pure SO<sub>3</sub> condenses into a trimer, which is often called γ-SO<sub>3</sub>. This molecular form is a colorless solid with a melting point of 16.8 °C. It adopts a cyclic structure described as [S(=O)<sub>2</sub>(μ-O)]<sub>3</sub>.<ref name="G&E">{{Greenwood&Earnshaw1st|pages=832–33}}.</ref> |
− | If SO<sub>3</sub> is condensed above 27 °C | + | If SO<sub>3</sub> is condensed above 27 °C in the presence of traces of water ([[amount fraction]] ''x'' = 10<sup>–5</sup>), then fibrous α-SO<sub>3</sub> forms, which has a melting point of 62.3 °C. Structurally, it is the [[polymer]] [S(=O)<sub>2</sub>(''μ''-O)]<sub>''n''</sub>, with each end of the polymer terminated by –OH groups. β-SO<sub>3</sub> is also fibrous, but of different molecular weight, also consisting of a hydroxyl-capped polymer and melting at 32.5 °C. Both the gamma- and the beta-forms are metastable, eventually converting to the more stable alpha-form if left standing for sufficient time in the presence of traces of water.<ref name="Merck">{{Merck9th|8775}}.</ref> |
Relative vapor pressures of solid SO<sub>3</sub> are alpha < beta < gamma at identical temperatures, indicative of their relative [[molecular weight]]s. Liquid sulfur trioxide has vapor pressure consistent with the gamma form. Thus heating a crystal of ''α''-SO<sub>3</sub> 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."<ref name="Merck"/> | Relative vapor pressures of solid SO<sub>3</sub> are alpha < beta < gamma at identical temperatures, indicative of their relative [[molecular weight]]s. Liquid sulfur trioxide has vapor pressure consistent with the gamma form. Thus heating a crystal of ''α''-SO<sub>3</sub> 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."<ref name="Merck"/> | ||
SO<sub>3</sub> is aggressively [[hygroscopic]]. In fact, the heat of hydration is sufficient that mixtures of SO<sub>3</sub> and wood or cotton can ignite. In such cases, SO<sub>3</sub> dehydrates these [[carbohydrate]]s.<ref name="Merck"/> | SO<sub>3</sub> is aggressively [[hygroscopic]]. In fact, the heat of hydration is sufficient that mixtures of SO<sub>3</sub> and wood or cotton can ignite. In such cases, SO<sub>3</sub> dehydrates these [[carbohydrate]]s.<ref name="Merck"/> | ||
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==References== | ==References== | ||
− | + | {{reflist}} | |
− | [[Category:Sulfur | + | [[Category:Sulfur compounds]] |
− | [[Category: | + | [[Category:Oxides]] |
[[Category:Acidic oxides]] | [[Category:Acidic oxides]] | ||
{{Imported from Wikipedia|name=Sulfur trioxide|id=303199049}} | {{Imported from Wikipedia|name=Sulfur trioxide|id=303199049}} |
Latest revision as of 17:44, 24 August 2009
Sulfur trioxide | |
---|---|
IUPAC name | Sulfur trioxide |
Other names | Sulfuric anhydride Sulfan |
Identifiers | |
InChI | InChI=1/O3S/c1-4(2)3 |
InChIKey | AKEJUJNQAAGONA-UHFFFAOYAX |
Standard InChI | InChI=1S/O3S/c1-4(2)3 |
Standard InChIKey | AKEJUJNQAAGONA-UHFFFAOYSA-N |
CAS number | [ ] |
EC number | |
UN number | 1829 |
RTECS | WT4830000 |
ChemSpider | |
Properties | |
Chemical formula | SO3 |
Molar mass | 80.06 g/mol |
Density | 1.92 g/cm3, liquid |
Melting point |
16.9 °C |
Boiling point |
45 °C |
Solubility in water | reacts violently |
Thermochemistry | |
Std enthalpy of formation ΔfH |
–397.77 kJ/mol |
Standard molar entropy S |
256.77 J K–1 mol–1 |
Hazards[1] | |
Material safety data sheet (MSDS) | ICSC |
EU index number | 016-019-00-2 |
GHS pictograms | |
GHS signal word | DANGER |
GHS hazard statements | H314, H335 |
Flash point | Non-flammable |
Related compounds | |
Other cations | Selenium trioxide Tellurium trioxide |
Other sulfur oxides | Sulfur monoxide Sulfur dioxide |
Other compounds | Sulfuric acid Disulfuric acid |
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) |
Sulfur trioxide (also spelled sulphur trioxide) is the chemical compound with the formula SO3. It is prepared industrially on massive scales 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:
- SO3 + H2O → H2SO4 (–88 kJ/mol)
The reaction occurs both rapidly and exothermically, too violently to be used in large-scale manufacturing. 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.[2]
Preparation
Sulfur trioxide can be prepared in the laboratory by the two-stage pyrolysis of sodium bisulfate. Sodium pyrosulfate is an intermediate product:
- Dehydration at 315°C:
- 2 NaHSO4 → Na2S2O7 + H2O
- Cracking at 460°C:
- Na2S2O7 → Na2SO4 + SO3
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.[3] 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.[4]
If SO3 is condensed above 27 °C in the presence of traces of water (amount fraction x = 10–5), then fibrous α-SO3 forms, which has a melting point of 62.3 °C. Structurally, it is the polymer [S(=O)2(μ-O)]n, with each end of the polymer terminated by –OH groups. β-SO3 is also fibrous, but of different molecular weight, also consisting of a hydroxyl-capped polymer and melting at 32.5 °C. Both the gamma- and the beta-forms are metastable, eventually converting to the more stable alpha-form if left standing for sufficient time in the presence of traces of water.[5]
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."[5]
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.[5]
References
- ↑ Index no. 016-019-00-2 of Annex VI, Part 3, to Regulation (EC) No 1272/2008 of the European Parliament and of the Council of 16 December 2008 on classification, labelling and packaging of substances and mixtures, amending and repealing Directives 67/548/EEC and 1999/45/EC, and amending Regulation (EC) No 1907/2006. OJEU L353, 31.12.2008, pp 1–1355 at p 400.
- ↑ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred Advanced Inorganic Chemistry, 6th ed.; Wiley-Interscience: New York, 1999. ISBN 0-471-19957-5
- ↑ Holleman, A. F.; Wiberg, E. Inorganic Chemistry; Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
- ↑ Greenwood, Norman N.; Earnshaw, A. Chemistry of the Elements; Pergamon: Oxford, 1984; pp 832–33. ISBN 0-08-022057-6.
- ↑ 5.0 5.1 5.2 The Merck Index: An Encyclopedia of Chemicals, Drugs, and Biologicals, 9th ed.; Merck, 8775.
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