What is Sulfur Trioxide?
Sulfur trioxide is described as a chemical compound. Sulfur trioxide formula is given as SO3. Sulfur trioxide is available in a number of modifications that varies in the form of molecular species and crystalline. It is colourless and forms liquid fumes in the air at ambient conditions. It is also a strong oxidising agent and a highly reactive substance, and it acts as a fire hazard. Thermodynamically, it is an unstable compound with respect to selenium dioxide.
The other names of sulfur trioxide can be given as sulfuric anhydride.
Structure of Sulphur Trioxide
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Properties of Sulfur Trioxide – SO3
Let us look at the important properties of sulfur trioxide given as follows:
Physical Properties of Sulfur Trioxide – SO3
Chemical Properties of Sulfur Trioxide – SO3
Sulfur trioxide compound reacts with water producing sulfuric acid. The chemical equation can be given below.
SO3 + H2O → H2SO4
Sulfur trioxide compound reacts with a base sodium hydroxide that produces sodium hydrogen phosphate. The chemical equation can be given below.
SO3 + NaOH → NaHSO4
Uses of Sulfur Trioxide – SO3
This can be used as a bleaching agent to remove the residual hydrogen peroxide, or as a digesting agent for pulp separation from lignin.
We can use it as a catalyst in the sulfur dioxide oxidation reaction to sulfur trioxide.
Strong inorganic acid mists that contain sulfuric acid, is used either in the industry or in the production of the commercial product.
It is also used as an important reagent in sulfonation reactions.
It can be used in the manufacturing of solar energy devices and photoelectric cells.
Molecular Structure and Bonding
SO3, in its gaseous form, is a D3h symmetry trigonal planar molecule, similarly predicted by VSEPR theory. So, it is said that SO3 belongs to the D3h point group.
In the aspect of electron-counting formalism, the sulfur atom contains an oxidation state of +6 and with a formal charge of 0. The Lewis structure holds one S=O double bond and two S–O dative bonds without utilising the d-orbitals.
The gaseous sulfur trioxide’s electrical dipole moment is given as zero. This is a consequence of the angle of 120° between the S-O bonds.
Let us look at the preparation of sulfur trioxide using various methods as follows:
In the Atmosphere
The direct oxidation of the sulfur dioxide to sulfur trioxide in air, and this reaction can be given as follows:
SO2 + 1⁄2O2 = SO3 ΔH=-198.4
The above reaction does take place, but this proceeds very slowly.
In the Laboratory
Sulfur trioxide is prepared in the chemical laboratory using the two-stage pyrolysis of the sodium bisulfate compound. The sodium pyrosulfate is given as an intermediate product:
At dehydration 315 °C, the chemical reaction is given as:
2 NaHSO4 → Na2S2O7 + H2O
Cracking at a temperature of 460 °C, the reaction can be given as:
Na2S2O7 → Na2SO4 + SO3
In contrast, KHSO4 compounds do not undergo a similar reaction.
It can also be prepared by adding the concentrated sulfuric acid to phosphorus pentoxide.
SO3 can be prepared industrially by the contact process. Sulfur dioxide, which in turn can be produced by the burning of iron pyrite (a sulfide ore of iron) or sulfur. After being purified by the electrostatic precipitation, the SO2 compound is then oxidised by atmospheric oxygen at a temperature between 400 and 600 °C over a catalyst. A typical catalyst contains vanadium pentoxide (V2O5) activated with the potassium oxide K2O on silica or kieselguhr support. Also, platinum works very well, but is much more expensive and is poisoned (which is rendered ineffective) much more easily by the impurities.
The majority sulfur trioxide compound, which is made in this way is converted into the sulfuric acid, but not by the direct addition of water, where it forms a fine mist, but by absorption in concentrated sulfuric acid and dilution with water of the formed oleum.
Once, it was industrially produced by heating the calcium sulfate with silica.
Sulfur trioxide is defined as an essential reagent in the sulfonation reactions. These processes afford dyes, pharmaceuticals, and detergents. Sulfur trioxide can be generated in situ from the sulfuric acid, or it can be used as a solution in the acid.
Sulfur trioxide, along with being a strong oxidising agent, will cause serious burns on both ingestion and inhalation because it is highly hygroscopic and corrosive in nature. SO3 compounds should be handled with extreme care because it reacts violently with water and also forms highly corrosive sulfuric acid. Also, it should be kept away from the organic material because of its strong dehydrating nature and its ability to react violently with such types of materials.
FAQs on Sulfur Trioxide - SO₃
1. How the Sulfur Trioxide and Water React?
Answer: To form sulfuric acid very vigorously, it needs a lot of heat. This is the reaction that can be used to form sulfuric acid by burning metal sulfides or sulfur to SO2, and then catalytically converting it to SO3, thereafter, reacting the SO3 with water.
However, we make SO3 gas in contact with the fairly concentrated sulfuric acid to make even more concentrated acid. Hence, we are reacting to a low concentration of water in H2SO4, and the reaction is less violent proportionately.
2. Why is Sulphur Trioxide Electrophilic in Nature?
Answer: Sulphur can be bonded to three oxygens, two times each. The sulphur present in the middle technically contains a formal charge of zero, whereas, the oxygens, which is bonded to are extremely electronegative (they are called electron hogs), and so the sulphur atom present in the middle contains a partial plus charge on it. Thus, it will accept the electrons to try to compensate.
3. Give the Charge of Sulfur.
Answer: Sulfur has an electron configuration of 2, 8, 6, which means that it could either gain 2 electrons and the charge becomes -2 or lose 6 electrons, and the charge will be +6
4. Explain if SO3 Acid or Base.
Answer: It can be said to be acidic. Because, Sulphur is a non-metal and we also know that generally, oxides of the non-metal are acidic. When the SO3 compound is dissolved in H2O, the solution turns blue litmus into the red by indicating the solution as acidic.