The sulfur oxide (VI) is an oily liquid, transparent and colorless. However, it may exhibit slight opacity and acquire an opaque white to light brown color. Under normal handling conditions, SO 3 and oleums have a high vapor pressure, SO 3 fumes react with ambient humidity to produce dense white clouds of sulfuric acid mist.
[ hide ]
- 1 Properties
- 1 Emission sources
- 2 Structure
- 1 Structure of the solid SO 3
- 3 Chemical properties
- 4 Source
- 5 External links
It is the product of the oxidation of sulfur oxide (IV) with oxygen in the presence of a catalyst such as vanadium pentoxide or platinum. It is produced on a large scale as a precursor to sulfuric acid .
It is a colorless and highly reactive gas that can easily condense into a liquid (Boiling point = 44.8 ºC). Under normal conditions, SO 3 is not found in the atmosphere because it reacts quickly with moisture, forming H 2 SO 4
Highly oxidizing, it is reduced by hydrogen and platinum.
The combustion of any substance containing sulfur will produce SO 2 and SO 3 . The relative amount of upper and lower oxides formed does not depend too much on the amount of oxygen present (unlike carbon oxides). Dioxide always forms in higher amounts under the conditions of any combustion. The amount of SO 3 produced depends on the reaction conditions, especially on the temperature, and ranges between 1% and 10% of the sulfur oxides produced.
Its origin is mainly in two causes:
- Homogeneous oxidation of SO 2
- Decomposition of sulfates present in fuels.
Of these, the homogeneous oxidation of SO 2 is the main one, so we can consider SO 3 as a secondary pollutant.
The gaseous form is a flat trigonal molecule of D3h symmetry, as predicted by the TREPEV theory. In the SO 3 molecule , the sulfur atom has an oxidation number of +6, with a formal charge of 0, and is surrounded by 6 pairs of electrons. From the perspective of molecular orbital theory, most of these electron pairs are nonbonding, typical behavior for hypervalent molecules.
Structure of the solid SO 3
The nature of the solid SO 3 is a surprisingly complex area because its structure changes due to traces of water.1 In the condensation of absolutely pure gas, sulfur oxide (IV) condenses into a trimer called γ-SO 3 . This molecular form is a colorless solid with a melting point of 16.8ºC and adopts a cyclic structure described as [S (= O) 2 (μ-O)] 3.2 If sulfur oxide (IV) condenses around 27ºC, α-SO3 is produced with a fibrous appearance, similar to asbestos (with which it has no chemical relationship). Structurally it is the polymer [S (= O) 2 (μ-O)] n. The chain ends at both ends with OH groups (hence α-SO3 is not really a form of SO 3 ). β-SO 3Like the alpha form, it has a fibrous appearance but has a different molecular weight, consisting of a hydroxyl polymer that melts at 32.5ºC. The gamma and beta forms are metastable, which eventually go into the stable alpha form if allowed long enough. This conversion is caused by traces of water.3 The relative pressures of solid sulfur (IV) oxide are alpha <beta <gamma, as indicated by their relative molecular weight. The liquid form of sulfur (IV) oxide matches the gamma form. So, heating an α-SO 3 crystal to its melting point causes a sudden increase in vapor pressure, which can be so powerful as to shatter the glass container in which it has been heated. This effect is known as the “alpha explosion” 3 SO 3it is aggressively hygroscopic. In fact, the heat of hydration from sulfur (IV) oxide mixtures and wood or cotton can burn, as SO 3 dehydrates carbohydrates.
In the presence of water it reacts violently giving rise to the formation of sulfuric acid, making it highly corrosive. SO3 is the acid anhydride of H 2 SO 4 , so the following reaction occurs:
SO 3 (l) + H 2 O (l) → H 2 SO 4 (l) (+88 kJ mol − 1)
The reaction is quick and exothermic. Around 340ºC, sulfuric acid, sulfur oxide (VI) and water coexist in significant equilibrium concentrations. Sulfur (IV) oxide also reacts with sulfur (II) chloride to produce thionyl chloride
SO 3 + SCl 2 → SOCl 2 + SO 2
Sulfur oxide (IV) can be prepared in the laboratory by two-stage pyrolysis from sodium hydrogen sulfate
2 NaHSO 4 → Na 2 S 2 O 7 + H 2 O at 315 ° C
Na 2 S 2 O 7 → Na 2 SO 4 + SO 3 at 460 ° C
This method works for other metal hydrogensulphates, the controlling factor being the stability of the intermediate salt of the pyrosulfate.
Industrially, sulfur oxide (VI) is obtained by contact process. Sulfur oxide (VI), generally obtained by burning sulfur or pyrite, is first purified by electrostatic precipitation. Sulfur oxide (VI) is oxidized purified in oxygen atmosphere at 400-600 over a catalyst of vanadium pentoxide V 2 O 5 activated oxide potassium K 2 O in silica support or kieselguhr. With platinum it also works very well but it is too expensive and is more easily contaminated by impurities.
Most of the sulfur oxide (VI) made in this way is converted into sulfuric acid , but not by the direct addition of water, with which it would form vapors; but by absorption in concentrated sulfuric acid and dilution with water of the oleum produced.