Sulfur trioxide

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For other uses, see SO3 (disambiguation).
This article is about the molecule SO
. For the ion SO2−
, see Sulfite.
Sulfur trioxide
Structure and dimensions of sulfur trioxide
Space-filling model of sulfur trioxide
Preferred IUPAC name
Sulfur trioxide
Systematic IUPAC name
Other names
Sulfuric anhydride
7446-11-9 YesY
ChEBI CHEBI:29384 YesY
ChemSpider 23080 YesY
EC Number 231-197-3
Jmol 3D model Interactive image
PubChem 24682
23035042 (monohydrate)
RTECS number WT4830000
UN number UN 1829
Molar mass 80.066 g/mol
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 give sulfuric acid
256.77 J K−1 mol−1
−395.7 kJ/mol
Vapor pressure {{{value}}}
Related compounds
Other cations
Selenium trioxide
Tellurium trioxide
Related sulfur oxides
Sulfur monoxide
Sulfur dioxide
Related compounds
Sulfuric acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YesYN ?)
Infobox references

Sulfur trioxide (alternative spelling, sulphur trioxide) is the chemical compound with the formula SO3. In the gaseous form, this species is a significant pollutant, being the primary agent in acid rain.[1] It is prepared on massive scales as a precursor to sulfuric acid.

Structure and bonding

Gaseous SO3 is a trigonal planar molecule of D3h symmetry, as predicted by VSEPR theory. SO3 belongs to the D3h point group.

In terms of electron-counting formalism, the sulfur atom has an oxidation state of +6 and a formal charge of +3. The Lewis structure consists of an S=O double bond and two S–O dative bonds without utilizing d-orbitals.[2]

The electrical dipole moment of gaseous sulfur trioxide is zero. This is a consequence of the 120° angle between the S-O bonds.

Chemical reactions

SO3 is the anhydride of H2SO4. Thus, the following reaction occurs:

SO3 (g) + H2O (l) → H2SO4 (aq) (ΔHf= −88 kJ mol−1)[citation needed]

The reaction occurs both rapidly and exothermically, too violently to be used in large-scale manufacturing. At or above 340 °C, sulfuric acid, sulfur trioxide, and water coexist in significant equilibrium concentrations.[citation needed]

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.[3]


Sulfur trioxide can be prepared in the laboratory by the two-stage pyrolysis of sodium bisulfate. Sodium pyrosulfate is an intermediate product:[4]

  1. Dehydration at 315 °C:
    2 NaHSO4 → Na2S2O7 + H2O
  2. Cracking at 460 °C:
    Na2S2O7 → Na2SO4 + SO3

This method will not work or will give a low yield for other metal bisulfates, particularly for potassium salt, because it easily gives away sulfur dioxide first, converting into dipotassium peroxymonosulfate, while the equivalent compound for sodium seems to be too unstable to form.[4]

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.[citation needed]

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.[citation needed]

Structure of solid SO3

Ball-and-stick model of the γ-SO3 molecule

The nature of solid SO3 is a surprisingly complex area because of structural changes caused by traces of water.[5] 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.[6]

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. Structurally, it is the polymer [S(=O)2(μ-O)]n. Each end of the polymer is terminated with OH groups. 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.[7]

Relative vapor pressures of solid SO3 are alpha < beta < gamma at identical temperatures, indicative of their relative molecular weights. Liquid sulfur trioxide has a 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".[7]

SO3 is aggressively hygroscopic. The heat of hydration is sufficient that mixtures of SO3 and wood or cotton can ignite. In such cases, SO3 dehydrates these carbohydrates.[7]


In process plant environment, SO3 gas is mixed into flue gas from combustion to give any particulates an electrostatic charge before they pass through electrostatic precipitators. The electrostatic precipitators will then trap the particulates, making cleaner process emission possible.[citation needed]


Sulfur trioxide will cause serious burns on both inhalation and ingestion because it is highly corrosive and hygroscopic in nature. SO3 should be handled with extreme care as it reacts with water violently and produces highly corrosive sulfuric acid.


See also


  1. Thomas Loerting, Klaus R. Liedl (2000). "Toward elimination of descrepancies between theory and experiment: The rate constant of the atmospheric conversion of SO3 to H2SO4". Proceedings of the National Academy of Sciences of the United States of America. 97 (16): 8874–8878. doi:10.1073/pnas.97.16.8874. 
  2. Terence P. Cunningham , David L. Cooper , Joseph Gerratt , Peter B. Karadakov and Mario Raimondi (1997). "Chemical bonding in oxofluorides of hypercoordinatesulfur". Journal of the Chemical Society, Faraday Transactions. 93 (13): 2247–2254. doi:10.1039/A700708F. 
  3. Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN 0-471-19957-5 
  4. 4.0 4.1
  5. Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5 
  6. Advanced Inorganic Chemistry by Cotton and Wilkinson, 2nd ed p543
  7. 7.0 7.1 7.2 Merck Index of Chemicals and Drugs, 9th ed. monograph 8775