SULFUR TRIOXIDE (SO3): STRUCTURE, PROPERTIES, RISKS, USES - CHEMISTRY - 2025

The sulfur trioxide is an inorganic compound formed by the union of a sulfur atom (S) and 3 oxygen atoms (O). Its molecular formula is SO ₃. At room temperature, SO ₃ is a liquid that releases gases into the air.

The structure of gaseous SO ₃ is flat and symmetric. All three oxygens are evenly located around sulfur. SO ₃ reacts violently with water. The reaction is exothermic, which means that heat is produced, in other words, it gets very hot.

Sulfur trioxide molecule SO ₃. Author: Benjah-bmm27. Source: Wikimedia Commons.

When liquid SO ₃ cools, it turns into a solid that can have three types of structure: alpha, beta, and gamma. The most stable is alpha, in the form of layers joined together to form a network.

Gaseous sulfur trioxide is used to prepare fuming sulfuric acid, also called oleum, because of its resemblance to oil or oily substances. Another of its important applications is in the sulfonation of organic compounds, that is, the addition of -SO ₃ - groups to them. Thus, useful chemicals such as detergents, colorants, pesticides, among many others, can be prepared.

SO ₃ is very dangerous, it can cause severe burns, damage to eyes and skin. Nor should it be inhaled or ingested as it can cause death from internal burns, in the mouth, esophagus, stomach, etc.

For these reasons, it must be handled with great caution. It must never come into contact with water or combustible materials such as wood, paper, fabrics, etc., as fires can occur. Neither should it be disposed of nor should it enter the sewers due to the danger of explosion.

The gaseous SO ₃ generated in industrial processes should not be released into the environment, as it is one of those responsible for the acid rain that has already damaged large areas of forests in the world.

Structure

The molecule of sulfur trioxide SO ₃ in the gaseous state has a triangular planar structure.

This means that both sulfur and the three oxygens are in the same plane. Furthermore, the distribution of oxygens and all electrons is symmetric.

Lewis resonance structures. Electrons are distributed evenly in SO ₃. Author: Marilú Stea.

In solid state three types of structure of SO _{3 are known}: alpha (α-SO ₃), beta (β-SO ₃) and gamma (γ-SO ₃).

The gamma γ-SO _{3 form} contains cyclic trimers, that is, three units of SO ₃ together forming a cyclic or ring-shaped molecule.

Gamma-type solid sulfur trioxide ring-shaped molecule. Author: Marilú Stea.

The beta β-SO ₃ phase has infinite helical chains of tetrahedra of composition SO ₄ linked together.

Structure of a chain of beta-type solid sulfur trioxide. Author: Marilú Stea.

The most stable form is alpha α-SO ₃, similar to beta but with a layered structure, with the chains joined to form a network.

Nomenclature

-Sulfur trioxide

-Sulfuric anhydride

-Sulfuric oxide

-SO ₃ gamma, γ-SO ₃

-SO ₃ beta, β-SO ₃

-SO ₃ alpha, α-SO ₃

Physical properties

Physical state

At room temperature (around 25 ºC) and atmospheric pressure, SO ₃ is a colorless liquid that emits smoke into the air.

When liquid SO ₃ is pure at 25 ºC it is a mixture of monomeric SO ₃ (a single molecule) and trimeric (3 joined molecules) of the formula S ₃ O ₉, also called SO ₃ gamma γ-SO ₃.

When lowering the temperature, if the SO ₃ is pure when it reaches 16.86 ºC, it solidifies or freezes to γ-SO ₃, also called “SO ₃ ice ”.

If it contains small amounts of moisture (even traces or extremely small amounts) SO ₃ polymerizes to the beta β-SO ₃ form which forms crystals with a silky shine.

Then more bonds are formed generating the alpha α-SO _{3 structure}, which is a needle-shaped crystalline solid that resembles asbestos or asbestos.

When alpha and beta merge they generate gamma.

Molecular weight

80.07 g / mol

Melting point

SO ₃ gamma = 16.86 ºC

Triple point

It is the temperature at which the three physical states are present: solid, liquid and gas. In the alpha form the triple point is at 62.2 ºC and in the beta it is at 32.5 ºC.

Heating the alpha form has a greater tendency to sublimate than to melt. Sublimate means to go from the solid to the gaseous state directly, without going through the liquid state.

Boiling point

All forms of SO ₃ boil at 44.8ºC.

Density

Liquid SO ₃ (gamma) has a density of 1.9225 g / cm ³ at 20 ºC.

Gaseous SO ₃ has a density of 2.76 relative to air (air = 1), which indicates that it is heavier than air.

Vapor pressure

SO ₃ alpha = 73 mm Hg at 25 ºC

SO ₃ beta = 344 mm Hg at 25 ºC

SO ₃ gamma = 433 mm Hg at 25 ºC

This means that the gamma form tends to evaporate more easily than beta and beta form than alpha.

Stability

The alpha form is the most stable structure, the others are metastable, that is, they are less stable.

Chemical properties

SO ₃ reacts vigorously with water to give sulfuric acid H ₂ SO ₄. When reacting, a lot of heat is produced so that water vapor is quickly released from the mixture.

When exposed to air, SO ₃ absorbs moisture quickly, emitting dense vapors.

It is a very strong dehydrating agent, this means that it removes water easily from other materials.

Sulfur in SO ₃ has an affinity for free electrons (that is, electrons that are not in a bond between two atoms) so it tends to form complexes with compounds that possess them, such as pyridine, trimethylamine or dioxane.

Complex between sulfur trioxide and pyridine. Benjah-bmm27. Source: Wikimedia Commons.

By forming complexes, sulfur “borrows” electrons from the other compound to fill its lack of them. Sulfur trioxide is still available in these complexes, which are used in chemical reactions to supply SO ₃.

It is a powerful sulfonating reagent for organic compounds, which means that it is used to easily add a group –SO ₃ - to molecules.

It reacts easily with the oxides of many metals to give sulfates of these metals.

It is corrosive to metals, animal and plant tissues.

SO ₃ is a difficult material to handle for several reasons: (1) its boiling point is relatively low, (2) it has a tendency to form solid polymers at temperatures below 30 ºC and (3) it has a high reactivity towards almost all organic substances and water.

May polymerize explosively if it does not contain a stabilizer and moisture is present. Dimethyl sulfate or boron oxide are used as stabilizers.

Obtaining

It is obtained by the reaction at 400 ºC between sulfur dioxide SO ₂ and molecular oxygen O ₂. However, the reaction is very slow and catalysts are required to increase the rate of the reaction.

2 SO ₂ + O ₂ ⇔ 2 SO ₃

Among the compounds that accelerate this reaction are platinum metal Pt, vanadium pentoxide V ₂ O ₅, ferric oxide Fe ₂ O ₃ and nitric oxide NO.

Applications

In the preparation of oleum

One of its main applications consists in the preparation of oleum or fuming sulfuric acid, so called because it emits vapors visible to the naked eye. To obtain it, SO ₃ is absorbed in concentrated sulfuric acid H ₂ SO ₄.

Oleum or fuming sulfuric acid. You can see the white smoke coming out of the bottle. W. Oelen. Source: Wikimedia Commons.

This is done in special stainless steel towers where the concentrated sulfuric acid (which is liquid) goes down and the gaseous SO ₃ goes up.

The liquid and the gas come into contact and come together, forming oleum which is an oily-looking liquid. It has a mixture of H ₂ SO ₄ and SO ₃, but it also has molecules of disulfuric acid H ₂ S ₂ O ₇ and trisulfuric acid H ₂ S ₃ O ₁₀.

In sulfonation chemical reactions

Sulfonation is a key process in large-scale industrial applications for the manufacture of detergents, surfactants, colorants, pesticides, and pharmaceuticals.

SO ₃ serves as a sulfonating agent to prepare sulfonated oils and alkyl-aryl-sulfonated detergents, among many other compounds. The following shows the sulfonation reaction of an aromatic compound:

ArH + SO ₃ → ArSO ₃ H

Sulfonation of benzene with SO ₃. Pedro8410. Source: Wikimedia Commons.

For the sulfonation reactions, oleum or SO ₃ can be used in the form of its complexes with pyridine or with trimethylamine, among others.

In the extraction of metals

SO ₃ gas has been used in mineral treatment. Simple oxides of metals can be converted to the much more soluble sulfates by treating them with SO ₃ at relatively low temperatures.

Sulfide minerals such as pyrite (iron sulfide), chalcosine (copper sulfide) and millerite (nickel sulfide) are the most economical sources of non-ferrous metals, so treatment with SO ₃ allows these metals to be easily obtained. and at low cost.

Iron, nickel and copper sulphides react with SO ₃ gas even at room temperature, forming the respective sulfates, which are very soluble and can be subjected to other processes to obtain the pure metal.

In various uses

SO _{3 is} used to prepare chlorosulfuric acid, also called chlorosulfonic acid HSO ₃ Cl.

Sulfur trioxide is a very powerful oxidant and is used in the manufacture of explosives.

Risks

To health

SO ₃ is a highly toxic compound by all routes, that is, inhalation, ingestion and contact with the skin.

Irritating and corroding mucous membranes. Causes skin and eye burns. Its vapors are very toxic when inhaled. Internal burns, shortness of breath, chest pain, and pulmonary edema occur.

Sulfur trioxide SO3 is very corrosive and dangerous. Author: OpenIcons. Source: Pixabay.

It is poisonous. Its ingestion generates severe burns of the mouth, esophagus and stomach. Furthermore, it is suspected of being a carcinogen.

From fire or explosion

It represents a fire hazard when in contact with materials of organic origin such as wood, fibers, paper, oil, cotton, among others, especially if they are wet.

There is also a risk if you come into contact with bases or reducing agents. It combines with water explosively, forming sulfuric acid.

Contact with metals can produce hydrogen gas H ₂ which is highly flammable.

Heating in glass jars should be avoided to prevent possible violent rupture of the container.

Environmental impact

SO ₃ is considered one of the major pollutants present in the earth's atmosphere. This is due to its role in the formation of aerosols and its contribution to acid rain (due to the formation of sulfuric acid H ₂ SO ₄).

Forest damaged by acid rain in the Czech Republic. Lovecz. Source: Wikimedia Commons.

SO ₃ is formed in the atmosphere by the oxidation of sulfur dioxide SO ₂. When SO _{3 is} formed, it reacts rapidly with water to form sulfuric acid H ₂ SO ₄. According to recent studies, there are other mechanisms for the transformation of SO ₃ in the atmosphere, but due to the large amount of water present in the atmosphere, it is still considered much more likely that SO ₃ turns mainly into H ₂ SO ₄.

SO ₃ gas or gaseous industrial waste containing it must not be discharged into the atmosphere because it is a dangerous pollutant. It is a highly reactive gas and, as mentioned above, in the presence of humidity in the air, SO ₃ turns into sulfuric acid H ₂ SO ₄. Therefore, in air, SO ₃ persists in the form of sulfuric acid, forming small droplets or aerosols.

If the sulfuric acid droplets enter the respiratory tract of humans or animals, they grow rapidly in size due to the moisture present there, so they have a chance to penetrate the lungs. One of the mechanisms by which the acid mist of H ₂ SO ₄ (that is, SO ₃) can produce strong toxicity is because it changes the extracellular and intracellular pH of living organisms (plants, animals and human beings).

According to some researchers, SO ₃ fog is the cause of the increase in asthmatics in an area of ​​Japan. The SO ₃ mist has a very corrosive effect towards metals, so that metal structures built by humans such as some bridges and buildings can be severely affected.

Liquid SO ₃ should not be disposed of in sewage drains or sewers. If spilled into sewers, it may create a fire or explosion hazard. If spilled by accident, do not direct a stream of water at the product. It should never be absorbed in sawdust or other combustible absorbent, as it can cause fires.

It must be absorbed in dry sand, dry earth or other totally dry inert absorbent. SO ₃ must not be released into the environment and must never be allowed to come into contact with it. It should be kept away from water sources because with this it produces sulfuric acid that is harmful to aquatic and terrestrial organisms.

References

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2. Muller, TL (2006). Sulfuric acid and sulfur trioxide. Kirk-Othmer Encyclopedia of Chemical Technology. Volume 23. Recovered from onlinelibrary.wiley.com.
3. US National Library of Medicine. (2019). Sulfur trioxide. Recovered from pubchem.ncbi.nlm.nih.gov.
4. Kikuchi, R. (2001). Environmental Management of Sulfur Trioxide Emission: Impact of SO ₃ on Human Health. Environmental Management (2001) 27: 837. Recovered from link.springer.com.
5. Cotton, F. Albert and Wilkinson, Geoffrey. (1980). Advanced Inorganic Chemistry. Fourth Edition. John Wiley & Sons.
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