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Sulfur Periodic Table Facts

Sulfur (S) or sulphur is the nonmetallic chemical element of Group 16 or VIA (Chalgens family) of the periodic table. Pure sulfur is a pale yellow color, tasteless, orderless, brittle solid, poor electrical conductor, and insoluble in water, uses widely for manufacturing sulfuric acid for the chemical industry. In 1809, the French chemists Louis-Joseph Gay-Lussac and Louis-Jacques Thenard discovered that the chemical elemental sulfur present in nature. It forms a large number of allotropes with puckered rings or long-chain sulfur atoms. In chemistry, except for gold and platinum, sulfur reacts with most of the elements to forms a large number of metallic and nonmetallic chemical compounds. Sulfur has four stable isotopes like 32S, 33S, 34S, and 36S.

Sulfur, nonmetals position on the periodic table chemical elements

The nonmetallic sulfur has chemical symbol S, atomic number 16, atomic weight melting point 115.21 °C, boiling point 444.6 °C, density 2.07 gm/cm3 (alpha), 1.96 gm/cm3 (beta), 1.92 gm/cm3 (gamma), and electron configuration 1s2 2s2 2p6 3s2 3p4. Sulfur shows a wide numbers of oxidation number in different types of chemical compounds like -2 (H2S), +4 (SO2), +6 (H2SO4).

Sulfur (S) or sulphur is a nonmetallic chemical element of Group 16 of the periodic table with properties and industrial uses

Chemical Properties

The outermost quantum orbital consists of an s2p4 electronic structure with a vacant d-orbital. In learning chemistry, the element can use their vacant d-orbitals for chemical bonding propose giving the valences 2, 4, and 6. The ionization energy is generally high for sulfur atom and decreases down the group of the period table. The normal trend of the metallic character increases with the increasing atomic number. Therefore, among Group 16 family, oxygen and sulfur is an insulator, selenium, and tellurium are semiconductors and polonium is a metal. The element is just two electrons short of the next noble gas configuration which can be achieved mostly by forming covalent bonding.

Abundance and Extraction

Sulfur, the tenth most abundant chemical element in the earth’s universe, is known from ancient history. Sulfur is found in many common minerals like gypsum (CaSO4, 2H2O), galena (PbS), pyrite (FeS2), cinnabar (HgS), and occurs as hydrogen sulfide in natural gas, petroleum, and as organosulfur compounds in coal. Today, nearly 25 percent of the element is produced from natural gas, petroleum, and ores containing sulfur materials. The native form is present in many places on the earth and sulfur occurs in many organic compounds like hair, wool, egg albumen, garlic, onion and mustered.

Sulfur is usually lifted from underground deposits by the Frasch process. Three concentric pipes are sunk to the sulfur layer of earth (150 to 200 meters below the surface), the element melted by superheated water and forced upwards by the air compressor to the extraction of the element. Recently, a large scale of sulfur extracted from the natural gases which contain 15 to 20 percent of hydrogen sulfide. The hydrogen sulfide absorbed in monoethanolamine follows by partial oxidation. The element is produced by self-reduction between H2S and SO2 over Fe2O3/Al2O3 chemical catalyst at 300 °C.

A million tonnes of elements found in the United States and Japan recovered from crude petroleum oil. Sulfate ores like gypsum treated with carbon to produce metal sulfide (CaS or BaS). The metal sulfide reacts with the acid to produce hydrogen sulfide which converts sulfur by the self-reduction process.


The allotrope contains a parked ring of 6 to 20 atoms. The allotrope practically differs in different types of packing units and varying repulsion between lone pairs of electrons on neighbor atoms. Cyclo-hexasulfur and cyclooctasufur contain non-planner puckered rings. The most common form of sulfur is the orthorhombic (Sα) with the S8 ring in crown formation. At 95.5 ºC, it slowly changes into monoclinic form (Sβ) with the S8 ring but in different packing. The monoclinic form is obtained by the crystallization of molten sulfur at about 100 °C, followed by rapid cooling to room temperature. A third crystalline solid variety, the γ-monoclinic form is obtained by chilling hot concentrated solutions of sulfur in CS2 or ethyl alcohol or by slowly cooling molten element from 150 °C.

The cyclo S6 species, also called Engel’s-sulfur, prepared by Engel in 1891 by adding hydrogen chloride to a saturated solution of sodium thiosulfate at 0 °C. Other cyclic species (S9, S10 up to S20) are obtained by reacting sulfur chloride with hydrogen. Chain species are collectively called catena-S available in very metastable forms like rubbery, plastic, laminar, fibrous, etc. Liquid sulfur (melting point 115 °C) consists almost entirely of S8 molecule. The viscosity of transparent yellow liquid first decreases with increasing temperature but after 159 °C, the viscosity rises steeply to reaches maximum at and thereafter fall smoothly to the boiling point. The changes due to the breaking of S8 rings to diradicals with unpaired electrons in terminal s-atoms.


Sulfur combines with the most of periodic table elements to forms most of the compounds in oxidation states -2 (sulfide, S-2), +4 (dioxide, SO2), and +6 (sulfate, SO4-2). It is the second element after carbon to exhibit catenation properties due to ring or chain formation. Hydride, sulfides, oxides, oxoacids, and salts are the most common examples of sulfur compounds.

Hydrogen sulfide or sulfurated hydrogen is the most familiar hydride of sulfur form by direct combination of the element with hydrogen. A large amount of hydrogen sulfide is obtained in the removal of sulfur from petroleum and widely uses as a reducing agent in chemical laboratories. Water formed by Group 16 element oxygen is a liquid due to the existence of hydrogen bonding but hydrogen sulfide is a gas with the characteristic smell of rotten eggs.

All the metals (except gold and platinum) react with S to form metal sulfide formed by ionic bonding and containing negatively charged sulfide ions (S-2). The sulfide of iron, nickel, copper, cobalt, and zinc are the imported ores of the respective metals. SO2 and SO3 are the main oxides prepared by heating the element with air or oxygen.

The oxides are soluble in water to give the acid solution of sulfurous and sulfuric acid. It is the major pollutant that causes air pollution, soil pollution, or acid rain in the earth’s environment. It reacts with fluorine to form sulfur hexafluoride (SF6) which is used as an insulator in the electric device. Sulfur is an important constituent of plants and animal bodies. Protein contains amino acids cysteine, cystine, and methionine involves many important metabolic reactions of the living organism.

Uses of Sulfur

  • Sulfur is widely (88 percent of total use) used in making sulfuric acid which is the topmost chemical in industry and state of the national economy.
  • Other important uses of sulfur include vulcanization of rubber, manufacturing of carbon disulfide (CS2), for making rayon, drugs, insecticides, and fungicides.
  • Sulfurous acid and SO2 are the essential bleaching agent in the sugar and paper industry.
  • Sulfuric acid obtained from sulfur or its compounds largely used for making phosphate and ammonium sulfate fertilizer, in making detergents, paints, pigments, rayon, petroleum products, sheet metal, storage batteries, explosive materials, and other different types of products.