The term “S molecule” refers to elemental sulfur, a nonmetallic chemical element with the symbol S and atomic number 16. It is the tenth most abundant element by mass in the universe and fifth on Earth. Sulfur is necessary for all living things and is used in a wide range of industrial processes.
Understanding the S Molecule: Structure and Forms
Elemental sulfur can form over 30 different solid allotropes—structural modifications of an element—which is more than any other element. This diversity arises from the tendency of sulfur atoms to link into rings or chains. The most common and stable form under normal conditions is octasulfur (S₈).
The S₈ molecule has a puckered, crown-shaped ring structure. The bond lengths between each sulfur atom are equal, measuring about 2.05 angstroms. This conformation is the basis for the most common forms of solid sulfur.
While the S₈ ring is the most stable, other sulfur rings exist, like the deeper yellow S₇ and orange-red S₆ molecules. Larger rings such as S₁₂ and S₁₈ have also been prepared and studied. Rapidly cooling molten sulfur produces an amorphous, non-crystalline form known as plastic sulfur, which may have a helical structure with eight atoms per turn.
Distinctive Properties of S Molecules
At room temperature, elemental sulfur is a bright yellow, crystalline solid. It is soft and odorless in its pure form, though some of its compounds, like hydrogen sulfide, are known for their strong smells. Sulfur is insoluble in water but will dissolve in nonpolar organic solvents like carbon disulfide. It has a melting point of 115.21 °C (239.38 °F) and a boiling point of 444.6 °C (832.3 °F).
When heated, sulfur melts into a blood-red liquid. Above 200°C (392°F), this liquid darkens and becomes more viscous due to the formation of polymer chains. When ignited, it burns with a blue flame, producing sulfur dioxide gas.
Sulfur is relatively unreactive at standard temperatures, but its reactivity increases with heat. It reacts with most metals to form sulfides and also combines with nonmetals. Its ability to accept electrons allows it to act as an oxidizing agent.
Natural Occurrence and Extraction of Sulfur
Elemental sulfur is found in its native form in regions with volcanic or geothermal activity. Significant underground deposits also occur, associated with salt domes and minerals like gypsum, particularly along the U.S. Gulf Coast. Sulfur is also widely distributed in minerals such as iron pyrites and galena.
Historically, the Frasch process was used to mine these deposits by pumping superheated water into the ground to melt the sulfur, which was then forced to the surface with compressed air. Today, most of the world’s sulfur is produced as a byproduct of processing fossil fuels. Natural gas and crude oil contain sulfur compounds that are removed during refining, and the Claus process is a common method used to convert these compounds into elemental sulfur.
Key Applications of S Molecules
The largest commercial use of elemental sulfur is for producing sulfuric acid (H₂SO₄), one of the most manufactured chemicals globally. A significant portion of this acid is used to create phosphate and sulfate fertilizers. It is also used in petroleum refining and to produce other chemicals and pigments.
In agriculture, elemental sulfur is used directly as a fungicide and insecticide. It is also applied to soil as a fertilizer to provide sulfur for plant growth and to help lower soil pH in alkaline conditions.
Another industrial application is the vulcanization of rubber. When natural rubber is heated with sulfur, the sulfur atoms form cross-links between the rubber molecules. This process makes the rubber harder, more durable, and more elastic, a development that was important for the tire industry.
Sulfur is also a component in a diverse array of products, including:
- Detergents
- Matches
- Fireworks
- Certain types of paper
In the pharmaceutical field, it is a component of sulfa drugs and is used in treatments for skin conditions.