Oct 02, 2019 Sulfur is an essential element for life. It's found in amino acids (cysteine and methionine) and proteins. Sulfur compounds are why onions make you cry, why asparagus gives urine a weird odor, why garlic has a distinctive aroma, and why rotten eggs smell so horrible. Tetrahedral Sulfur, 67-9004, is represented by Tetrahedral Phosphorus, 67-6924, when used with four Long Connector Link, 67-7021, resulting in a covalent radius CPK Atomic Models, Sulfur Atom, Tetrahedral. Sulfur forms over 30 solid allotropes, more than any other element. Besides S 8, several other rings are known. Removing one atom from the crown gives S 7, which is more of a deep yellow than the S 8. HPLC analysis of 'elemental sulfur' reveals an equilibrium mixture of mainly S 8, but with S 7 and small amounts of S 6. Atomic Number – Protons, Electrons and Neutrons in Sulfur. Sulfur is a chemical element with atomic number 16 which means there are 16 protons in its nucleus.Total number of protons in the nucleus is called the atomic number of the atom and is given the symbol Z. 3, 1, -1, -1/2 Sulfur = element 16 The first quantum number tells you which energy level electron 16 is in, it is in level 3 The second number tells you which sublevel the electron is in - it is in the 3p sublevel. We use s=0, p=1, d=2, f=3 The third quantum number tells you which orbital of the sublevel an electron is in. The three orbitals of the 3p sublevel will each fill with an up spin.

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Sulfur (S), also spelled sulphur, nonmetallic chemical element belonging to the oxygen group (Group 16 [VIa] of the periodic table), one of the most reactive of the elements. Pure sulfur is a tasteless, odourless, brittle solid that is pale yellow in colour, a poor conductor of electricity, and insoluble in water. It reacts with all metals except gold and platinum, forming sulfides; it also forms compounds with several nonmetallic elements. Millions of tons of sulfur are produced each year, mostly for the manufacture of sulfuric acid, which is widely used in industry.

In cosmic abundance, sulfur ranks ninth among the elements, accounting for only one atom of every 20,000–30,000. Sulfur occurs in the uncombined state as well as in combination with other elements in rocks and minerals that are widely distributed, although it is classified among the minor constituents of Earth’s crust, in which its proportion is estimated to be between 0.03 and 0.06 percent. On the basis of the finding that certain meteorites contain about 12 percent sulfur, it has been suggested that deeper layers of Earth contain a much larger proportion. Seawater contains about 0.09 percent sulfur in the form of sulfate. In underground deposits of very pure sulfur that are present in domelike geologic structures, the sulfur is believed to have been formed by the action of bacteria upon the mineralanhydrite, in which sulfur is combined with oxygen and calcium. Deposits of sulfur in volcanic regions probably originated from gaseous hydrogen sulfide generated below the surface of Earth and transformed into sulfur by reaction with the oxygen in the air.

History

The history of sulfur is part of antiquity. The name itself probably found its way into Latin from the language of the Oscans, an ancient people who inhabited the region including Vesuvius, where sulfur deposits are widespread. Prehistoric humans used sulfur as a pigment for cave painting; one of the first recorded instances of the art of medication is in the use of sulfur as a tonic.

The combustion of sulfur had a role in Egyptian religious ceremonials as early as 4,000 years ago. “Fire and brimstone” references in the Bible are related to sulfur, suggesting that “hell’s fires” are fuelled by sulfur. The beginnings of practical and industrial uses of sulfur are credited to the Egyptians, who used sulfur dioxide for bleaching cotton as early as 1600 bce. Greek mythology includes sulfur chemistry: Homer tells of Odysseus’ use of sulfur dioxide to fumigate a chamber in which he had slain his wife’s suitors. The use of sulfur in explosives and fire displays dates to about 500 bce in China, and flame-producing agents used in warfare (Greek fire) were prepared with sulfur in the Middle Ages. Pliny the Elder in 50 ce reported a number of individual uses of sulfur and ironically was himself killed, in all probability by sulfur fumes, at the time of the great Vesuvius eruption (79 ce). Sulfur was regarded by the alchemists as the principle of combustibility. Antoine Lavoisier recognized it as an element in 1777, although it was considered by some to be a compound of hydrogen and oxygen; its elemental nature was established by the French chemists Joseph Gay-Lussac and Louis Thenard.

Natural occurrence and distribution

Many important metal ores are compounds of sulfur, either sulfides or sulfates. Some important examples are galena (lead sulfide, PbS), blende (zinc sulfide, ZnS), pyrite (iron disulfide, FeS₂), chalcopyrite (copper iron sulfide, CuFeS₂), gypsum (calcium sulfate dihydrate, CaSO₄∙2H₂O) and barite (barium sulfate, BaSO₄). The sulfide ores are valued chiefly for their metal content, although a process developed in the 18th century for making sulfuric acid utilized sulfur dioxide obtained by burning pyrite. Coal, petroleum, and natural gas contain sulfur compounds.

Allotropy

Sulfur Atomic Number

In sulfur, allotropy arises from two sources: (1) the different modes of bonding atoms into a single molecule and (2) packing of polyatomic sulfur molecules into different crystalline and amorphous forms. Some 30 allotropic forms of sulfur have been reported, but some of these probably represent mixtures. Only eight of the 30 seem to be unique; five contain rings of sulfur atoms and the others contain chains.

In the rhombohedral allotrope, designated ρ-sulfur, the molecules are composed of rings of six sulfur atoms. This form is prepared by treating sodium thiosulfate with cold, concentrated hydrochloric acid, extracting the residue with toluene, and evaporating the solution to give hexagonal crystals. ρ-sulfur is unstable, eventually reverting to orthorhombic sulfur (α-sulfur).

A second general allotropic class of sulfur is that of the eight-membered ring molecules, three crystalline forms of which have been well characterized. One is the orthorhombic (often improperly called rhombic) form, α-sulfur. It is stable at temperatures below 96 °C (204.8 °F). Another of the crystalline S₈ ring allotropes is the monoclinic or β-form, in which two of the axes of the crystal are perpendicular, but the third forms an oblique angle with the first two. There are still some uncertainties concerning its structure; this modification is stable from 96 °C to the melting point, 118.9 °C (246 °F). A second monoclinic cyclooctasulfur allotrope is the γ-form, unstable at all temperatures, quickly transforming to α-sulfur.

An orthorhombic modification, S₁₂ ring molecules, and still another unstable S₁₀ ring allotrope are reported. The latter reverts to polymeric sulfur and S₈. At temperatures above 96 °C (204.8 °F), the α-allotrope changes into the β-allotrope. If enough time is allowed for this transition to occur completely, further heating causes melting to occur at 118.9 °C (246 °F); but if the α-form is heated so rapidly that the transformation to β-form does not have time to occur, the α-form melts at 112.8 °C (235 °F).

Just above its melting point, sulfur is a yellow, transparent, mobile liquid. Upon further heating, the viscosity of the liquid decreases gradually to a minimum at about 157 °C (314.6 °F), but then rapidly increases, reaching a maximum value at about 187 °C (368.6 °F); between this temperature and the boiling point of 444.6 °C (832.3 °F), the viscosity decreases. The colour also changes, deepening from yellow through dark red, and, finally, to black at about 250 °C (482 °F). The variations in both colour and viscosity are considered to result from changes in the molecular structure. A decrease in viscosity as temperature increases is typical of liquids, but the increase in the viscosity of sulfur above 157 °C probably is caused by rupturing of the eight-membered rings of sulfur atoms to form reactive S₈ units that join together in long chains containing many thousands of atoms. The liquid then assumes the high viscosity characteristic of such structures. At a sufficiently high temperature, all of the cyclic molecules are broken, and the length of the chains reaches a maximum. Beyond that temperature, the chains break down into small fragments. Upon vaporization, cyclic molecules (S₈ and S₆) are formed again; at about 900 °C (1,652 °F), S₂ is the predominant form; finally, monatomic sulfur is formed at temperatures above 1,800 °C (3,272 °F).

Study Notes

The chemistry of sulfur‑containing organic compounds is often omitted from introductory organic chemistry courses. However, we have included a short section on these compounds, not for the sake of increasing the amount of material to be digested, but because much of the chemistry of these substances can be predicted from a knowledge of their oxygen‑containing analogues.

A thiol is a compound which contains an $ce{-}$SH functional group. The $ce{-}$SH group itself is called a mercapto group. A disulfide is a compound containing an $ce{-}$S$ce{-}$S$ce{-}$ linkage.

(Organic) sulfides have the structure R$ce{-}$S$ce{-}$R′, and are therefore the sulfur analogues of ethers. The nomenclature of sulfides can be easily understood if one understands the nomenclature of the corresponding ethers. Notice that the term “thio” is also used in inorganic chemistry. For example, SO₄²⁻ is the sulfate ion; while S₂O₃²⁻, in which one of the oxygen atoms of a sulfate ion has been replaced by a sulfur atom, is called thiosulfate.

Thiolate anions, RS$ce{-}$, are analogous to alkoxy anions, RO$ce{-}$. Thiolate anions are better nucleophiles than are alkoxy anions.

If you have trouble understanding why trialkylsulfonium ions are formed, think of them as being somewhat similar to the hydronium ions that are formed by protonating water:

Later we shall see examples of tetraalkylammonium ions, R₄N⁺, which again may be regarded as being similar to hydronium ions.

Sulfoxides and sulfones are obtained by oxidizing organic sulfides. You need not memorize the methods used to carry out these oxidations.

Table 18.1, below, provides a quick comparison of oxygen‑containing and sulfur‑containing organic compounds.

Sulfur Atomic Size

Note that when we name thiols, we include the “e” of the alkane name. Thus, CH₃CH₂SH is called “ethanethiol,” not “ethanthiol.”