Sulfur is one of the elements known since ancient times, used long before the concept of a chemical element existed. Ancient civilizations (Egyptians, Greeks, Romans, Chinese) used native sulfur found near volcanoes for religious, medicinal, and industrial purposes. The ancient Greeks used sulfur to bleach fabrics and as a fumigant, while the Romans used it in their metallurgical processes. In medieval alchemy, sulfur was considered one of the three fundamental principles, along with mercury and salt. In 1777, Antoine Lavoisier (1743-1794) definitively demonstrated that sulfur is a chemical element and not a compound, including it in his list of elements in his Traité élémentaire de chimie (1789). The name sulfur comes from the Latin sulfur, probably of Sanskrit origin.
Sulfur (symbol S, atomic number 16) is a non-metal of group 16 (formerly group VIA, the chalcogen family) of the periodic table. Its atom has 16 protons, 16 electrons, and usually 16 neutrons in its most abundant isotope (\(\,^{32}\mathrm{S}\)). Four stable isotopes exist: sulfur-32 (\(\,^{32}\mathrm{S}\)), sulfur-33 (\(\,^{33}\mathrm{S}\)), sulfur-34 (\(\,^{34}\mathrm{S}\)), and sulfur-36 (\(\,^{36}\mathrm{S}\)).
At room temperature, sulfur is a lemon-yellow crystalline solid, brittle, and odorless in its elemental form (the characteristic odor comes from its compounds like H₂S). Density ≈ 2.07 g/cm³. Sulfur exhibits remarkable polymorphism with many allotropic forms. The main forms are orthorhombic sulfur α (S₈, stable at room temperature, cyclic octahedra), melting point: 388.36 K (115.21 °C), and monoclinic sulfur β (stable above 95.3 °C). Boiling point of sulfur: 717.8 K (444.6 °C). Molten sulfur has extraordinary viscous properties that vary with temperature.
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Sulfur-32 — \(\,^{32}\mathrm{S}\,\) | 16 | 16 | 31.972071 u | ≈ 94.99 % | Stable | By far the most abundant isotope of natural sulfur. |
| Sulfur-33 — \(\,^{33}\mathrm{S}\) | 16 | 17 | 32.971459 u | ≈ 0.75 % | Stable | Used in isotopic geochemistry to trace ancient biological processes. |
| Sulfur-34 — \(\,^{34}\mathrm{S}\) | 16 | 18 | 33.967867 u | ≈ 4.25 % | Stable | Important in geochemistry for studying biogeochemical cycles. |
| Sulfur-36 — \(\,^{36}\mathrm{S}\) | 16 | 20 | 35.967081 u | ≈ 0.01 % | Stable | Rare isotope; used as a tracer in environmental research. |
| Sulfur-35 — \(\,^{35}\mathrm{S}\) | 16 | 19 | 34.969032 u | Not natural | 87.37 days | Radioactive β\(^-\) decaying into chlorine-35. Used as a tracer in biology and medicine. |
| Other isotopes — \(\,^{26}\mathrm{S}\) to \(\,^{49}\mathrm{S}\) | 16 | 10 — 33 | — (variable) | Not natural | Milliseconds to seconds | Unstable isotopes produced artificially; experimental nuclear physics. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Sulfur has 16 electrons distributed across three electron shells. Its full electronic configuration is: 1s² 2s² 2p⁶ 3s² 3p⁴, or simplified as: [Ne] 3s² 3p⁴. This configuration can also be written as: K(2) L(8) M(6).
K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 8 electrons distributed as 2s² 2p⁶. This shell is also complete, forming a noble gas configuration (neon).
M Shell (n=3): contains 6 electrons distributed as 3s² 3p⁴. The 3s orbitals are complete, while the 3p orbitals contain only 4 out of 6 possible electrons. Thus, 2 electrons are missing to saturate this outer shell.
The 6 electrons in the outer shell (3s² 3p⁴) are the valence electrons of sulfur. This configuration explains its chemical properties:
By gaining 2 electrons, sulfur forms the S²⁻ ion (oxidation state -2), a common state in metallic sulfides, thus adopting the configuration of argon [Ar].
By losing or sharing electrons, sulfur can exhibit various positive oxidation states: +2, +4, and +6, the latter being observed in sulfuric acid H₂SO₄ and sulfates.
The oxidation state 0 corresponds to elemental sulfur, which exists in various allotropic forms, the most stable being orthorhombic sulfur S₈ (a ring of 8 atoms).
The electronic configuration of sulfur, with 6 electrons in its valence shell, classifies it among the chalcogens (elements of group 16). This structure gives it characteristic properties: great chemical versatility with many possible oxidation states (-2, 0, +2, +4, +6), the ability to form chains and rings through S-S bonding, and the ability to form both ionic and covalent bonds. Sulfur can accept 2 electrons to achieve the stability of a noble gas, forming the sulfide ion S²⁻ present in many minerals. It can also share its electrons in covalent bonds, forming a wide variety of compounds. Its chemical versatility makes sulfur an essential element in industrial chemistry, particularly for the production of sulfuric acid (the most produced chemical compound in the world), rubber vulcanization, and in many important organosulfur compounds in biology.
Sulfur is moderately reactive at room temperature but very reactive at high temperatures. It combines with almost all chemical elements except noble gases and nitrogen. Sulfur burns in air with a characteristic blue flame, producing irritating and suffocating sulfur dioxide (SO₂). It forms compounds in several oxidation states: -II (sulfides and thiols), +IV (SO₂, sulfites) and +VI (SO₃, sulfates, sulfuric acid H₂SO₄). Sulfur reacts with metals to form sulfides, with hydrogen to form hydrogen sulfide (H₂S, a toxic gas with a rotten egg smell), and with oxygen to form various oxides. The chemistry of sulfur is extremely rich, including polysulfur chains, organosulfur compounds, and disulfide bonds essential in biochemistry.
Sulfur is the sixth essential element for life (C, H, N, O, P, S). It is present in two essential amino acids: cysteine and methionine, which make up the proteins of all living beings. Disulfide bridges (S-S bonds) between cysteine residues are crucial for the three-dimensional structure and stability of proteins. Sulfur is also present in several vital coenzymes such as coenzyme A, biotin (vitamin B₇), and lipoic acid. It plays an important role in cellular detoxification via glutathione, a powerful antioxidant. Some bacteria use sulfur in their energy metabolism (sulfate-reducing and sulfur-oxidizing bacteria), playing a major role in the biogeochemical sulfur cycle. Volcanic hot springs rich in sulfur host unique extremophile microbial ecosystems.
Sulfur is the tenth most abundant element in the universe and is found naturally in several forms. In the Earth's crust, it represents about 0.035% of the mass. Native sulfur (elemental) is found in volcanic regions and near hot springs. It also exists in many minerals: metal sulfides (pyrite FeS₂, galena PbS, sphalerite ZnS), sulfates (gypsum CaSO₄·2H₂O, barite BaSO₄), and in fossil fuels (oil, coal, natural gas). The sulfur cycle involves complex biological, geological, and atmospheric processes. Volcanic emissions, decomposition of organic matter, and industrial activity (combustion of fossil fuels) release sulfur dioxide into the atmosphere, contributing to acid rain when it transforms into sulfuric acid.
Sulfur is a relatively abundant element in the universe, produced during nucleosynthesis in massive stars by the fusion of oxygen and silicon. Supernovae disperse significant amounts of sulfur into the interstellar medium. Sulfur has been detected in many celestial objects: stellar atmospheres, nebulae, comets, meteorites, and planetary atmospheres. Jupiter's moon Io exhibits active volcanism dominated by sulfur, giving its surface a characteristic yellow-orange color. Sulfur geysers on Io eject sulfur dioxide hundreds of kilometers high. Venus has clouds containing sulfuric acid in its atmosphere. The detection of sulfur compounds such as H₂S in the atmospheres of exoplanets could constitute an indirect biosignature.
N.B.:
Molten sulfur exhibits extraordinary and counterintuitive viscous behavior. When solid sulfur is heated, it melts around 115 °C into a light yellow, fluid liquid. But when heated further above 160 °C, the liquid suddenly becomes extremely viscous, almost solid, changing from a water-like consistency to that of thick honey, then molasses. This phenomenon results from the breaking of S₈ rings, which form entangled polymeric chains. Upon further heating above 200 °C, the viscosity gradually decreases again. This unique behavior makes molten sulfur a fascinating system for studying phase transitions and polymerization.
