Uses of Sulphur Dioxide

The chemical substance with the SO₂ is sulphur dioxide (IUPAC-recommended spelling) or sulphur dioxide (traditional Commonwealth English). It’s a poisonous gas that gives off the odour of burned matches. It is created by copper extraction and the combustion of sulphur-bearing fossil fuels and is released naturally by volcanic activity. The odour of sulphur dioxide is the same as that of nitric acid.
Sulphur dioxide is a precursor of sulphuric acid since it is turned to sulphur trioxide and subsequently to oleum, which is then converted to sulphuric acid. Sulphur dioxide is created when sulphur reacts with oxygen for this purpose. The contact process is the method of turning sulphur dioxide into sulphuric acid. Every year, several billion kg are manufactured for this purpose.
There are a variety of stable sulphur oxides, but sulphur dioxide and sulphur trioxide are the most frequent. Sulphur dioxide is a sulphur oxide that is often found. It’s a nonflammable, colourless gas with a distinct odour. It has a melting point of -72.5 C and a density of 2.8 kg/m³. The liquid form is an excellent solvent because organic molecules are more soluble in SO₂ than water. The most common usage of SO₂ is to make SO₃. SO₂ is produced by burning sulphur directly and roasting metal sulphides.

The Chemical Named Sulphur

Sulphur is a chemical element with the atomic number 16 and the chemical symbol “S” on the periodic chart. Following strontium, sulphur is the seventeenth most abundant element, accounting for 0.0384 percent of the Earth’s crust. Elemental sulphur, organo-sulphur compounds in oil and coal, H₂S(g) in natural gas, and mineral sulphites and sulphates are all examples of sulphur. The Frasch process (described below) extracts this element, which involves drawing liquid sulphur to the surface with superheated water and compressed air. The principal sites that supply large amounts of elemental sulphur include offshore areas, Texas and Louisiana. However, decreasing H₂S, widely found in oil and natural gas, can also yield elemental sulphur. On the other hand, sulphur is mainly utilised to make SO₂(g) and H₂SO₄. Sulphur has been since referred to as brimstone in the Hebrew texts, and it was defined as an element by Lavoisier in 1777. Pure sulphur has a pale yellow colour and is tasteless and odourless. Sulphur samples commonly encountered in the lab have a distinct odour. The tenth most abundant element in this vast universe is sulphur.

Physical Properties of Sulphur

Sulphur belongs to the oxygen family and has an atomic weight of 32.066 grams per mole. It has a specific heat of 0.706 Jg^{-1 o}C^-1 and is a nonmetal. The electronegativity is 2.58, and the electron affinity is 200 kJ mol-1 (unitless). In vast numbers of tiny orthorhombic crystals, sulphur is often a light-yellow, opaque, and brittle solid. Sulphur is twice as dense as water, but it is also insoluble in water. On the other hand, sulphur is very soluble in carbon disulfide and very mildly soluble in various other solvents. Due to carbonaceous impurities, sulphur can also change colour and blacken when heated. Sulphur is substantially darkened by even 0.05 percent carbonaceous materials. Most sulphur is extracted directly from underground deposits by injecting super-heated water and piping out molten sulphur (sulphur melts at 112o C). Sulphur has the most allotropes of any element. Although the S8 ring is the most frequent, there are six additional allotropes with 20 sulphur atoms per ring. O=O is more powerful than S=S, whereas O–O is less potent than S–S. As a result, single bonds are preferred in sulphur compounds, and catenation is easier than oxygen compounds. The explanation for the weaker S=S double bonds appears to be related to the atom’s size: it’s more difficult for the two atoms to come together at a close enough distance for the p orbitals to overlap, resulting in a weak bond. Looking down the periodic table confirms this: The bond enthalpy of Se=Se is even lower, at 272kJ/mol.

Uses of Sulphur Dioxide

1. Used as a preservative:

Sulphur dioxide can prevent microbe-caused deterioration and increase the shelf life of food. In addition to its antibacterial characteristics, sulphur dioxide is an antioxidant that inhibits browning caused by oxidation in foods, particularly dried fruits and vegetables – this aids in the preservation of the items’ look and colour. Light-coloured fruits, such as dried apricots, can quickly darken without sulphur dioxide exposure to air. Because of its antibacterial characteristics and capacity to prevent oxidation, sulphur dioxide is occasionally employed as a preservative in dried apricots, dried figs, and other types of dried fruits. It acts as a preservative, preserving the freshness of the fruit and preventing it from spoiling. Sulfites, bisulfites, and metabisulfites, as well as their neutral and acid salts, are employed in meals in gaseous or liquid form. In addition to its antibacterial properties, SO₂ is used in foods as an antioxidant, bleaching agent, colour fixative, inhibitor of enzymatic discolourations, and nonenzymatic browning inhibitor. Fruits and vegetables, fruit juices and concentrates, syrups, wines and jams, and to a lesser extent, prawns, fish, minced meats, sausages, and mushrooms are all regularly preserved using SO₂.

2. Used in biochemical and biomedical rules:

Sulphur contributes significantly to nature’s chemical diversity, and its special qualities permit vital biological interactions that no other element can. All living creatures employ natural sulphur compounds, which are divided into many kingdoms according to their roles. Marine organisms are one of the most important sources of natural sulphur products since most inorganic sulphur is metabolised in ocean environments where this element is abundant. Sulphur may also be integrated into organic molecules by terrestrial organisms like plants and bacteria, resulting in basic metabolites (like methionine and cysteine) and more complex chemical compounds with a range of biological activities. Sulphur dioxide, or its conjugate base bisulfite, is produced by both sulphate-reducing organisms and sulphur-oxidising bacteria. Furthermore, these metabolic markers were closely linked to the severity of pulmonary arterial hypertension. Sulphur dioxide is a powerful anti-inflammatory, antioxidant, and cytoprotective substance generated naturally. It lowers blood pressure and reduces hypertension blood vessel remodelling, especially in time thickening, in hypertensive patients. It’s also involved in lipid metabolism.

3. Used as reducing agent and reagent:

Sulphur dioxide can be employed as a reducing or oxidising agent. Sulphur dioxide is commonly used as a reducing agent to make living easier. Sulphur dioxide can decolourise things in the presence of water. It’s particularly effective as a decreasing bleach for papers and delicate fabrics like clothing. Because of its excellent reductant, it is utilised as a refrigerant in refrigerators. Sulphur dioxide is used to detoxify chlorinated wastewater before being released in municipal wastewater treatment. In the presence of water, sulphur dioxide may discolour objects. It works well as a reducing bleach for papers and other delicate materials such as textiles.
SO₂ + H₂O ⇌ HSO₃⁻ + H⁺  

Conclusion

Sulphur dioxide has a wide range of properties that make it a compound of high importance and necessity through all means and mediums. Whenever the study of oxides enters the domain of teaching, sulphur oxide becomes the most exquisite and necessary compound. Given the numerous applications and its wide range of benefits, it finds greater importance in various strata of chemistry.