Sulfur trioxide

The molecule SO₃istrigonal 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).^[7] 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.^[1] The electrical dipole 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 3 polymorphs are known, with conversion between them being dependent on traces of water.^[9]

Absolutely pure SO₃ freezes at 16.8 °C to give the γ-SO₃ form, which adopts the cyclic trimer configuration [S(=O)₂(μ-O)]₃.^[10][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. Structurally, it is the polymer [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 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".^[11]

Sulfur trioxide undergoes many reactions.^[1]

SO₃ is the anhydride of H₂SO₄. Thus, it is susceptible to hydration:

SO₃ + H₂O → H₂SO₄ (Δ_fH = −200 kJ/mol)^[12]

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.^[11]

Akin to the behavior of H₂O, hydrogen fluoride adds to give fluorosulfuric acid:

SO₃ + HF → FSO₃H

SO₃ reacts with dinitrogen pentoxide to give the nitronium salt of pyrosulfate:

2 SO₃ + N₂O₅ → [NO₂]₂S₂O₇

Sulfur trioxide is an oxidant. It oxidizes sulfur dichloridetothionyl chloride.

SO₃ + SCl₂ → SOCl₂ + SO₂

SO₃ is a strong Lewis acid readily forming adducts with Lewis bases.^[13] With pyridine, it gives the sulfur trioxide pyridine complex. Related adducts form from dioxane and trimethylamine.

Sulfur trioxide is a potent sulfonating agent, i.e. it adds SO₃ groups to substrates. Often the substrates are organic, as in aromatic 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:

SO₂ + 1⁄2O₂ → SO₃ (ΔH = −198.4 kJ/mol)

Industrially SO₃ is made by the contact process. Sulfur dioxide is produced by the burning of sulfuroriron 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 oxideK₂O on kieselguhrorsilica 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 into sulfuric acid.

Sulfur trioxide can be prepared in the laboratory by the two-stage pyrolysisofsodium bisulfate. Sodium pyrosulfate is an intermediate product:^[17]

1. Dehydration at 315 °C:

   2 NaHSO₄ → Na₂S₂O₇ + H₂O

2. Cracking at 460 °C:

   Na₂S₂O₇ → Na₂SO₄ + SO₃

The latter occurs at much lower temperatures (45–60 °C) in the presence of catalytic H₂SO₄.^[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:

1. 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

2. Pyrolysis of anhydrous tin(IV) sulfate at 150 °C - 200 °C:

   Sn(SO₄)₂ → SnO₂ + 2 SO₃

The advantage of this method over the sodium bisulfate one is that it requires much lower temperatures and can be done using normal borosilicate laboratory glassware without 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.^[19]

Sulfur trioxide is a reagent in sulfonation reactions. These processes afford 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.^[20]

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

• Hypervalent molecule
• Sulfur trioxide pyridine complex

• NIST Standard Reference Database
• ChemSub Online