SULFUR: HISTORY, PROPERTIES, STRUCTURE, OBTAINING, USES - CHEMISTRY - 2025

The sulfur is a nonmetallic element leads, under oxygen, the group of chalcogens of the periodic table. It is specifically located in group 16 with period 3, and is represented by the chemical symbol S. Of its natural isotopes, ³² S is by far the most abundant (about 94% of all sulfur atoms).

It is one of the most abundant elements on Earth, making up about 3% of its total mass. In other words, if all the sulfur on the planet were taken, two yellow moons could be built; there would be three satellites instead of one. It can adopt various oxidation states (+2, -2, +4 and +6), so its salts are numerous and enrich the earth's crust and core.

Sulfur crystals. Source: Pixabay.

Sulfur is synonymous with yellow, bad smells and hell. The main reason for its bad odors is due to its derived compounds; especially sodas and organic ones. Of the rest, its minerals are solid and have colors that include yellow, gray, black and white (among others).

It is one of the elements that most presents a large number of allotropes. It can be found as small, discrete molecules of S ₂ or S ₃; as rings or cycles, being the orthorhombic and monoclinic sulfur S ₈ the most stable and abundant of all; and as helical chains.

It is not only found in the earth's crust in the form of minerals, but also in the biological matrices of our bodies. For example, it is in the amino acids cystine, cysteine ​​and methionine, in iron proteins, keratin and in some vitamins. It is also present in garlic, grapefruit, onions, cabbage, broccoli and cauliflower.

Chemically it is a soft element, and in the absence of oxygen it forms sulfurous minerals and sulfates. It burns with a bluish flame and may appear as an amorphous or crystalline solid.

Despite being essential for the synthesis of sulfuric acid, a highly corrosive substance, and having unpleasant odors, it is actually a benign element. Sulfur can be stored in any space without major precautions, as long as fires are avoided.

History of sulfur

In the Bible

Sulfur is one of the oldest elements in the history of mankind; so much so that its discovery is uncertain and it is not known which of the ancient civilizations used it for the first time (4000 years before Christ). In the very pages of the Bible, he can be found accompanying hellfire and hell.

The supposed smell of sulfur from hell is believed to have to do with volcanic eruptions. Its first discoverer must surely have come across mines of this element such as dust lands or yellow crystals in the vicinity of a volcano.

Antiquity

This yellowish solid soon demonstrated remarkable healing effects. For example, the Egyptians used sulfur to treat inflammation of the eyelids. It also relieved scabies and acne, an application that can be seen today in sulfur soaps and other dermatological items.

The Romans used this element in their rituals, as a fumigant and bleach. When it burns, it releases SO ₂, a gas that flooded the rooms, mixing with the humidity and providing an antibacterial medium capable of killing insects.

The Romans, like the Greeks, discovered the high combustibility of sulfur, which is why it became synonymous with fire. The color of its bluish flames must have illuminated the Roman circuses. It is believed that the Greeks, for their part, used this element to create incendiary weapons.

The Chinese, for their part, learned that by mixing sulfur with saltpeter (KNO ₃) and coal, they created the material black powder that set a historic turn, and that aroused great demands and interest in this mineral in the nations of the time.

Modern times

As if gunpowder wasn't reason enough to covet sulfur, sulfuric acid and its industrial applications soon emerged. And with the rod of sulfuric acid, the amount of wealth or prosperity of a country was measured in relation to its levels of consumption of this compound.

It was not until 1789 that the brilliant chemist Antoine Lavoisier was able to recognize sulfur and classify it as an element. Then in 1823 the German chemist Eilhard Mitscherlich discovered that sulfur can predominantly crystallize in two ways: rhombohedral and monoclinic.

The history of sulfur followed the same course of its compounds and applications. With the enormous industrial importance of sulfuric acid, it was accompanied by the vulcanization of rubbers, the synthesis of penicillin, the exploitation of mines, the refining of crude oil rich in sulfur, the nutrition of soils, etc.

Properties

Physical appearance

Brittle solid in powder or crystal form. Its color is dull lemon yellow, it is tasteless and has no odor.

Liquid appearance

Liquid sulfur is unique in that its initial yellow color turns reddish and intensifies and darkens if subjected to high temperatures. When it burns, it emits bright blue flames.

Molar mass

32 g / mol.

Melting point

115.21 ° C.

Boiling point

445 ° C.

ignition point

160 ° C.

Auto ignition temperature

232 ° C.

Density

2.1 g / mL. However, other allotropes may be less dense.

Molar heat capacity

22.75 J / mol K

Covalent radius

105 ± 3 pm.

Electronegativity

2.58 on the Pauling scale.

Polarity

SS bonds are apolar because both sulfur atoms have the same electronegativity. This makes all of its allotropes, cyclic or chain-shaped, nonpolar; and therefore, its interactions with water are inefficient and it cannot be solubilized in it.

However, sulfur can be dissolved in nonpolar solvents such as carbon disulfide, CS ₂, and aromatics (benzene, toluene, xylene, etc.).

Ion

Sulfur can form various ions, usually anions. The best known of all is sulfur, S ^2-. S ^2- is characterized by being bulky and has a soft Lewis base.

Because it is a soft base, the theory states that it will tend to form compounds with soft acids; such as transition metal cations, including Fe ²⁺, Pb ²⁺ and Cu ²⁺.

Structure and electronic configuration

The sulfur crown

S8 molecule, the most stable and abundant allotrope of sulfur. Source: Benjah-bmm27.

Sulfur can occur in a wide variety of allotropes; and these in turn have crystalline structures that are modified under different pressures and / or temperatures. Therefore, sulfur is an element rich in allotropes and polymorphs, and the study of its solid structures represents an endless source of theoretical-experimental work.

Why such structural complexity? To begin with, the covalent bonds in sulfur (SS) are very strong, being surpassed only by those of carbon, CC, and by that of hydrogen, HH.

Sulfur, unlike carbon, does not tend to form tetrahedra but boomerangs; that with their angles fold and ring to stabilize the sulfur chains. The best-known ring of all, also representing the most stable allotrope of sulfur, is S ₈, the “sulfur crown” (top image).

Note that all the SS links in the S ₈ look like individual boomerangs, resulting in a ring with pleats and not flat at all. These S ₈ crowns interact through London forces, orienting themselves in such a way that they create structural patterns that define an orthorhombic crystal; called S ₈ α (S-α, or simply orthorhombic sulfur).

Polymorphs

The sulfur crown is one of the many allotropes for this element. S ₈ α is a polymorph of this crown. There are two others (among the most important) called S ₈ β and S ₈ γ (S-β and S-γ, respectively). Both polymorphs crystallize into monoclinic structures, with S ₈ γ being denser (gamma sulfur).

All three are yellow solids. But how do you get each polymorph separately?

S ₈ β is prepared by heating S ₈ α to 93 ° C, then allowing its slow cooling to slow its transition back to the orthorhombic phase (the α). And S ₈ γ, on the other hand, is obtained when S ₈ α melts at 150 ° C, again allowing it to cool slowly; it is the densest of the sulfur crown polymorphs.

Other cyclic allotropes

The crown S ₈ is not the only cyclic allotrope. There are others such as S ₄, S ₅ (analogous to cyclopentane), S ₆ (represented by a hexagon like cyclohexane), S ₇, S ₉, and S _10-20; the latter means that there may be rings or cycles containing from ten to twenty sulfur atoms.

Each of them represents different cyclic allotropes of sulfur; and in turn, to emphasize it, they have varieties of polymorphs or polymorphic structures that depend on pressure and temperature.

For example, S ₇ has up to four known polymorphs: α, β, γ, and δ. The members or crowns of higher molecular masses are products of organic synthesis and do not predominate in nature.

Sulfur chains

Sulfur chain. Source: OpenStax

As more sulfur atoms are incorporated into the structure, their tendency to ring decreases and the sulfur chains remain open and adopt helical conformations (as if they were spirals or screws).

And thus, another voluminous family of sulfur allotropes emerges that does not consist of rings or cycles but of chains (like the one in the image above).

When these SS chains line up in parallel in the crystal, they trap impurities and end up defining a fibrous solid called fibrous sulfur, or S-ψ. If between these parallel chains there are covalent bonds that interconnect them (as happens with the vulcanization of rubber), we have laminar sulfur.

When sulfur S ₈ melts, a yellowish liquid phase is obtained that can turn dark if the temperature is increased. This is because SS bonds are broken, and therefore a thermal depolymerization process occurs.

This liquid when cooled shows plastic and then glassy characteristics; that is, a vitreous and amorphous sulfur (S-χ) is obtained. Its composition consists of both rings and sulfur chains.

And when a mixture of the fibrous and laminar allotrope is obtained from the amorphous sulfur, Crystex is produced, a commercial product used for the vulcanization of rubber.

Small allotropes

Although they are left last, they are no less important (or interesting) than the allotropes of higher molecular masses. The S ₂ and S ₃ molecules are the sulfurized versions of O ₂ and O ₃. In the first, two sulfur atoms are joined with a double bond, S = S, and in the second there are three atoms with resonance structures, S = SS.

Both S ₂ and S ₃ are gaseous. S ₃ shows a cherry red color. Both have enough bibliographic material to each cover an individual article.

Electronic configuration

The electron configuration for the sulfur atom is:

3s ² 3p ⁴

It can gain two electrons to complete its valence octet, and thus have an oxidation state of -2. Likewise, it can lose electrons, starting with two in its 3p orbitals, its oxidation state being +2; if you lose two more electrons, with their 3p orbitals empty, your oxidation state will be +4; and if you lose all the electrons, it will be +6.

Obtaining

Mineralogical

Sulfur is part of many minerals. Among them are pyrite (FeS ₂), galena (PbS), covellite (CuS), and other sulfate and sulfide minerals. By processing them, not only the metals can be extracted, but also the sulfur after a series of reductive reactions.

It can also be obtained in a pure way in volcanic vents, where as the temperature rises it melts and spills downhill; And if it catches fire, it will look like bluish lava at night. Through arduous labor, and strenuous physical labor, sulfur can be harvested just as it was done quite often in Sicily.

Sulfur can also be found in underground mines, which are made to pump superheated water to melt it and move it to the surface. This obtaining process is known as the Frasch Process, currently little used.

Oil

Today most of the sulfur comes from the oil industry, as its organic compounds are part of the composition of crude oil and its refined derivatives.

If a crude or refined product is rich in sulfur and undergoes hydrodesulfurization, it will release large amounts of H ₂ S (stinky gas that smells like rotten eggs):

RSR + 2 H ₂ → 2 RH + H ₂ S

The H ₂ S is then chemically treated in the Clauss process, summarized with the following chemical equations:

3 O ₂ + 2 H ₂ S → 2 SO ₂ + 2 H ₂ O

SO ₂ + 2 H ₂ S → 3 S + 2 H ₂ O

Applications

Some of the uses for sulfur are mentioned below and in a general way:

- It is an essential element for both plants and animals. It is even present in two amino acids: cysteine ​​and methionine.

- It is the raw material for sulfuric acid, a compound involved in the preparation of countless commercial products.

- In the pharmaceutical industry it is used for the synthesis of sulfur derivatives, penicillin being the best known of the examples.

- Allows the vulcanization of rubbers by interconnecting the polymeric chains with SS bonds.

- Its yellow color and its mixtures with other metals make it desirable in the pigment industry.

- Mixed with an inorganic matrix, such as sand and rocks, concrete and sulfur asphalt are prepared to replace bitumen.

Risks and precautions

Sulfur by itself is a harmless, non-toxic substance, and it also poses no potential risks, unless it reacts to form other compounds. Its sulfate salts are not dangerous and can be handled without major precautions. This is not the case, however, with its gaseous derivatives: SO ₂ and H ₂ S, both highly toxic.

If it is in the liquid phase, it can cause serious burns. If it is swallowed in large quantities it can trigger the production of H ₂ S in the intestines. Otherwise, it does not represent any risk for those who chew it.

In general terms, sulfur is a safe element that does not require too many precautions, except to keep it away from fire and strong oxidizing agents.

References

1. Shiver & Atkins. (2008). Inorganic chemistry. (Fourth edition). Mc Graw Hill.
2. Laura Crapanzano. (2006). Polymorphism of sulfur: Structural and Dynamical Aspects. Physics.Université Joseph-Fourier - Grenoble I. English. fftel-00204149f
3. Wikipedia. (2019). Allotropes of sulfur. Recovered from: en.wikipedia.org
4. Meyer Beat. (1976). Elemental sulfur. Chemical Reviews, Vol. 76, No. 3.
5. Dr. Doug Stewart. (2019). Sulfur Element Facts. Chemicool. Recovered from: chemicool.com
6. Donald W. Davis and Randall A. Detro. (2015). Sulfur History. Georgia Gulf Sulfur Corporation. Recovered from: georgiagulfsulfur.com
7. Helmenstine, Anne Marie, Ph.D. (January 11, 2019). 10 Interesting Sulfur Facts. Recovered from: thoughtco.com
8. Boone, C.; Bond, C.; Hallman, A.; Jenkins, J. (2017). Sulfur General Fact Sheet; National Pesticide Information Center, Oregon State University Extension Services. npic.orst.edu