Sulfur Trioxide

Sulfur trioxide (alternative spelling sulphur trioxide) is the chemical compound with the formula SO₃. It has been described as "unquestionably the most conomicallyimportant sulfur oxide". It is prepared on an industrial scale as a precursor to sulfuric 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 to acid rain.

Molecular structure and bonding

Monomer

The molecule SO₃ is trigonal planar. As predicted by VSEPR theory, its structure belongs to the D_3h point group. The sulfur atom has an oxidation state of +6 and may be assigned a formal charge value as low as 0 (if all three sulfur-oxygen bonds are assumed to be double bonds) or as high as +2 (if the Octet Rule is assumed). When the formal charge is non-zero, the S-O bonding is assumed to be delocalized. In any case the three S-O bond lengths are equal to one another, at 1.42 Å. The electrical dipole moment of gaseous sulfur trioxide is zero.

Trimer

Both liquid and gaseous SO₃ exists in an equilibrium between the monomer and the cyclic trimer. The nature of solid SO₃ is complex and at least 3 polymorphs are known, with conversion between them being dependent on traces of water.

Absolutely pure SO₃ freezes at 16.8 °C to give the ''γ''-SO₃ form, which adopts the cyclic trimer configuration (=O)₂(''μ''-O)sub>3.

Polymer

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 (=O)₂(''μ''-O)sub>''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.''Merck Index of Chemicals and Drugs'', 9th ed. monograph 8775

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".

Chemical reactions

Sulfur trioxide undergoes many reactions.

Hydration and hydrofluorination

SO₃ is the anhydride of H₂SO₄. Thus, it is susceptible to hydration:
:SO₃ + H₂O → H₂SO₄(Δ_f''H'' = −200 kJ/ mol)

Gaseous sulfur trioxide fumes profusely even in a relatively dry atmosphere owing to formation of a sulfuric acid mist.
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.

Akin to the behavior of H₂O, hydrogen fluoride adds to give fluorosulfuric acid:
:SO₃ + HF → FSO₃H

Deoxygenation

SO₃ reacts with dinitrogen pentoxide to give the nitronium salt of pyrosulfate:
:2 SO₃ + N₂O₅ → O₂sub>2S₂O₇

Oxidant

Sulfur trioxide is an oxidant. It oxidizes sulfur dichloride to thionyl chloride.
:SO₃ + SCl₂ → SOCl₂ + SO₂

Lewis acid

SO₃ is a strong Lewis acid readily forming adducts with Lewis bases. With pyridine, it gives the sulfur trioxide pyridine complex. Related adducts form from dioxane and trimethylamine.

Sulfonating agent

Sulfur trioxide is a potent sulfonating agent, i.e. it adds SO₃ groups to substrates. Often the substrates are organic, as in aromatic sulfonation. For activated substrates, Lewis base adducts of sulfur trioxide are effective sulfonating agents.

Preparation

The direct oxidation of sulfur dioxide to sulfur trioxide in air proceeds very slowly:

: 2 SO₂ + O₂ → 2 SO₃(Δ''H'' = −198.4 kJ/mol)

Industrial

Industrially SO₃ is made by the contact process. Sulfur dioxide is produced by the burning of sulfur or iron pyrite (a sulfide ore of iron). After being purified by electrostatic precipitation, the SO₂ is then oxidised by atmospheric oxygen at between 400 and 600 °C over a catalyst. A typical catalyst consists 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 sulfur trioxide made in this way is converted into sulfuric acid.

Laboratory

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

The latter occurs at much lower temperatures (45–60 °C) in the presence of catalytic H₂SO₄. In contrast, KHSO₄ undergoes the same reactions at a higher temperature.

Another two step method involving a salt pyrolysis starts with concentrated sulfuric acid and anhydrous tin tetrachloride:

# Reaction between tin tetrachloride and sulfuric acid in a 1:2 molar mixture at near reflux (114 °C):
#: SnCl₄ + 2 H₂SO₄ → Sn(SO₄)₂ + 4 HCl
# Pyrolysis of anhydrous tin(IV) sulfate at 150 °C - 200 °C:
#:Sn(SO₄)₂ → SnO₂ + 2 SO₃

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 dehydrating sulfuric acid with phosphorus pentoxide.

Applications

Sulfur trioxide is a reagent in sulfonation reactions. Dimethyl sulfate is produced commercially by the reaction of dimethyl ether with sulfur trioxide:
:

Sulfate esters are used as detergents, dyes, and pharmaceuticals. Sulfur trioxide is generated in 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.

Safety

Along with being an oxidizing agent, sulfur trioxide is highly corrosive. It reacts violently with water to produce highly corrosive sulfuric acid.

See also

* Hypervalent molecule
* Sulfur trioxide pyridine complex

References

Sources

NIST Standard Reference Database

{{DEFAULTSORT:Sulfur Trioxide
Sulfur oxides
Acid anhydrides
Acidic oxides
Interchalcogens
Hypervalent molecules
Sulfur(VI) compounds