Sulfur trioxide

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.

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Structure and bonding

Gaseous SO₃ is a trigonal planar molecule of D_3h 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(l) + H₂O(l) → H₂SO₄(l) (-88 kJ mol⁻¹)

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.^[2] 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^[3].

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^[4].

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."^[4]

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.^[4]

Sources

• WebElements
• NIST Standard Reference Database
• European Chemicals Bureau

See also

• Hypervalent molecule

References

1. ^ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th Edn.) New York:Wiley-Interscience. ISBN 0-471-19957-5.
2. ^ Holleman, A. F.; Wiberg, E. "Inorganic Chemistry" Academic Press: San Diego, 2001. ISBN 0-12-352651-5.
3. ^ Advanced Inorganic Chemistry by Cotton and Wilkinson, 2nd ed p543
4. ^ ^{a b c} Merck Index of Chemicals and Drugs, 9th ed. monograph 8775