Octasulfur
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| Systematic IUPAC name | |||
| Other names
Octasulfur
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| 2973 | |||
| MeSH | Cyclooctasulfur | ||
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CompTox Dashboard (EPA)
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| Properties | |||
| S8 | |||
| Molar mass | 256.48 g·mol−1 | ||
| Appearance | Vivid, yellow, translucent crystals | ||
| Density | 2.07 g/cm3 | ||
| Melting point | 119 °C; 246 °F; 392 K | ||
| Boiling point | 444.6 °C; 832.4 °F; 717.8 K | ||
| log P | 6.117 | ||
| Thermochemistry | |||
Std molar
entropy (S⦵298) |
32 J·mol−1·K−1[3] | ||
Std enthalpy of
formation (ΔfH⦵298) |
0 kJ/mol[3] | ||
| Related compounds | |||
Related compounds
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Hexathiane | ||
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Octasulfur is an inorganic substance with the chemical formula S8. It is an odourless and tasteless yellow solid, and is a major industrial chemical. It is the most common allotrope of sulfur and occurs widely in nature.[4]
Nomenclature
[edit]The name octasulfur is the most commonly used for this chemical. It is systematically named cyclo-octasulfur (which is the preferred IUPAC name) and cyclooctasulfane. It is also the final member of the thiocane heterocylic series, where every carbon atom is substituted with a sulfur atom, thus this sulfur allotrope is systematically named octathiocane as well.
Structure
[edit]The chemical consists of rings of 8 sulfur atoms. It adopts a crown conformation with D4d point group symmetry. The S–S bond lengths are equal, at about 2.05 Å. Octasulfur crystallizes in three distinct polymorphs: rhombohedral, and two monoclinic forms, of which only two are stable at standard conditions. The rhombohedral crystal form is the accepted standard state. The remaining polymorph is only stable between 96 and 115 °C at 100 kPa. Octasulfur forms several allotropes: α-sulfur, β-sulfur, γ-sulfur, and λ-sulfur.
λ-Sulfur is the liquid form of octasulfur, from which γ-sulfur can be crystallised by quenching. If λ-sulfur is crystallised slowly, it will revert to β-sulfur. Since it must have been heated over 115 °C, neither crystallised β-sulfur or γ-sulfur will be pure. The only known method of obtaining pure γ-sulfur is by crystallising from solution.
Octasulfur easily forms large crystals, which are typically yellow and are somewhat translucent.
Production and reactions
[edit]It is the main (99%) component of elemental sulfur, which is recovered from volcanic sources and is a major product of the Claus process, associated with petroleum refineries. In space, sulfur can be formed by energetic processing of solid hydrogen sulfide,[5] which explains the recent detection of octasulfur in the return samples from the carbonaceous asteroid Ryugu.[6]
See also
[edit]References
[edit]- ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. p. 49. Electronic version.
- ^ "cyclooctasulfur (CHEBI:29385)". Chemical Entities of Biological Interest. UK: European Bioinformatics Institute. Main.
- ^ a b Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 978-0-618-94690-7.
- ^ Steudel, R., "Homocyclic Sulfur Molecules", Topics Curr. Chem. 1982, 102, 149.
- ^ Herath, Ashanie; Mcanally, Mason; Turner, Andrew M.; Wang, Jia; Marks, Joshua H.; Foretenberry, Ryan C.; Garcia-Alvarez, Jorge C.; Gozem, Samer; Kaiser, Ralf I. (1 July 2025). "Missing interstellar sulfur in inventories of polysulfanes and molecular octasulfur crowns". Nature Communications. doi:10.1038/s41467-025-61259-2. PMC 12215914.
- ^ Takano, Yoshinori (10 July 2024). "Primordial aqueous alteration recorded in water-soluble organic molecules from the carbonaceous asteroid (162173) Ryugu". Nature Communications. doi:10.1038/s41467-024-49237-6. PMC 11237059.
External links
[edit]
Media related to Octasulfur at Wikimedia Commons
Octasulfur
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Nomenclature and Structure
Nomenclature
Octasulfur, denoted as S₈, is the most stable and prevalent molecular form of elemental sulfur under standard conditions. It is commonly referred to by the name "octasulfur," which directly indicates its composition of eight sulfur atoms, or "crown sulfur," alluding to the puckered ring arrangement of its atoms.[3][4] The systematic IUPAC name for this compound is octathiocane, a term derived from the heterocyclic nomenclature where "thiocane" denotes an eight-membered ring containing sulfur, extended here to an all-sulfur cycle. Alternative designations include cyclooctasulfane and cyclo-octasulfur, emphasizing its cyclic nature as the final member in the series of polysulfane rings.[5] Historically, elemental sulfur, including its octasulfur allotrope, has been known since prehistoric times under names like "brimstone," from Old English "brennen stān" meaning "burning stone," due to its ready ignition. In the 17th to 19th centuries, a purified, finely powdered form obtained by sublimation was termed "flowers of sulfur" (or "flores sulphuris" in Latin), which primarily comprises S₈ molecules and was used in early chemical and medicinal applications. The shift to modern systematic nomenclature, such as octasulfur and its IUPAC equivalents, emerged in the early 20th century alongside advancements in structural chemistry and the establishment of IUPAC guidelines for inorganic compounds.[6][7]Molecular Structure
Octasulfur, with the molecular formula S, consists of eight sulfur atoms arranged in a cyclic structure. This molecule adopts a puckered crown conformation, characterized by D point group symmetry, where the ring is non-planar to minimize steric repulsion between lone pairs on adjacent atoms.[8][9] In this arrangement, the S–S bond lengths average 2.05 Å. The intraring bond angles are consistently around 107.8°, reflecting the tetrahedral-like geometry influenced by the valence shell electron pair repulsion. These structural parameters have been precisely determined through single-crystal X-ray diffraction studies at low temperatures, confirming the uniformity of the ring in isolated or solution phases.[10][11] The bonding in octasulfur involves homonuclear single σ-bonds between sulfur atoms, formed by overlap of p-orbitals, with each sulfur atom retaining two lone pairs in its valence shell. This configuration aligns with sulfur's ground-state electron arrangement ([Ne] 3s 3p), where two electrons per atom are shared in bonds, leaving four as non-bonding pairs that occupy hybrid orbitals. The absence of π-bonding contributes to the molecule's reactivity and stability.[12] Compared to smaller cyclic sulfur allotropes such as S and S, the S ring exhibits greater thermodynamic stability, as evidenced by its higher second-order energy difference and lower fragmentation energy, making it the predominant form under ambient conditions. While S (D symmetry) and S can be synthesized and isolated, they are less favored due to increased ring strain and higher energy states relative to the optimally sized eight-membered crown.[13]Polymorphs and Allotropes
Octasulfur exhibits three primary crystalline polymorphs, each consisting of packed S8 rings in distinct arrangements. The α-polymorph, orthorhombic with space group Fddd, is the most stable at room temperature and features 128 sulfur atoms per unit cell, corresponding to 16 S8 molecules arranged in a body-centered lattice.[14] Its unit cell parameters are approximately a = 10.44 Å, b = 12.85 Å, and c = 24.37 Å, resulting in a density of 2.07 g/cm³. The β-polymorph, monoclinic with space group P21/c, forms at higher temperatures and packs 48 sulfur atoms (6 S8 molecules) per unit cell in a less dense configuration, with parameters a ≈ 9.43 Å, b ≈ 11.70 Å, c ≈ 10.07 Å, and β ≈ 107.9°.[15] The γ-polymorph, also monoclinic (P21/c), achieves the highest density among these forms at 2.19 g/cm³ through tighter S8 ring packing, with unit cell parameters a ≈ 9.89 Å, b ≈ 10.82 Å, c ≈ 9.75 Å, and β ≈ 100.9°; it is often considered a standard reference state due to its preparation from precipitated sulfur. In all polymorphs, the S8 rings adopt a crown-like conformation, with intermolecular interactions dominated by van der Waals forces between rings. These octasulfur polymorphs relate to broader sulfur allotropy, where λ-sulfur (also known as fibrous or φ-sulfur) represents a polymeric form with helical catena chains rather than rings, formed by rapid quenching of molten sulfur.[16] Catena-sulfur encompasses various chain-like structures, including long linear polymers that emerge above the λ-transition temperature of approximately 159°C in the melt, contrasting with the cyclic S8 dominance in the solid polymorphs. Octasulfur polymorphs remain the prevalent solid forms up to 95.5°C, beyond which β predominates until melting. Phase transitions among these polymorphs are temperature-dependent. The α-to-β conversion occurs reversibly at 95.5°C, allowing β to revert to α upon cooling below this threshold.[17] In contrast, the γ-polymorph undergoes an irreversible transition to α upon heating or prolonged exposure at room temperature.[18] Recent analyses of samples from the Ryugu asteroid, returned by the Hayabusa2 mission, have detected S8 octasulfur alongside other sulfur species, providing evidence of extraterrestrial occurrence in carbonaceous materials and insights into solar system sulfur distribution as of 2024.[19]Physical and Thermodynamic Properties
Physical Characteristics
Octasulfur, in its stable α-form, manifests as a vivid yellow, translucent crystalline solid composed of orthorhombic crystals. This characteristic appearance arises from the packing of S₈ rings in a puckered crown conformation, giving it a distinctive bright hue under ambient conditions.[2] The material exhibits a density of 2.07 g/cm³ at 20°C, reflecting its compact molecular structure. Its melting point is 112.8°C for the α-form, transitioning to a viscous yellow liquid, while the boiling point reaches 444.6°C under standard pressure.[20] These thermal transitions are influenced by polymorphic variations, though the α-form predominates at room temperature.[21] Octasulfur demonstrates negligible solubility in water, rendering it hydrophobic and unreactive in aqueous environments. In contrast, it dissolves readily in nonpolar organic solvents, such as carbon disulfide (39 g/100 mL at 25°C), benzene (approximately 1.2 g/100 mL at 25°C), and toluene (approximately 2 g/100 mL at 25°C).[22] The compound possesses a faint odor, often described as mild or nearly odorless in pure form, and exhibits low volatility due to its minimal vapor pressure at room temperature.[2] Optically, octasulfur in the α-form has a refractive index of 1.957, contributing to its translucency and high light dispersion in crystalline samples.[20]Thermodynamic Data
Octasulfur, in its standard rhombic form (α-S₈), serves as the reference state for sulfur's thermodynamic properties, with the standard enthalpy of formation ΔH_f° and Gibbs free energy of formation ΔG_f° both defined as 0 kJ/mol at 298 K and 1 bar, underscoring its thermodynamic stability as the elemental standard.[23] The standard molar entropy S° is 32.1 J/mol·K at 298 K, reflecting the ordered molecular structure in the solid phase.[23] The molar heat capacity at constant pressure C_p for the solid is 22.74 J/mol·K at 298 K, indicating moderate thermal response typical of covalent network solids.[24] Phase transitions of octasulfur involve characteristic energy changes that govern its behavior under varying conditions. The enthalpy of fusion is 1.727 kJ/mol at the melting point of 112.8°C, representing the energy required to disrupt the solid lattice into the liquid phase.[25] The enthalpy of vaporization is approximately 45 kJ/mol at the normal boiling point of 444.6°C, highlighting the significant intermolecular forces in the liquid that must be overcome for gas formation.[26] The phase diagram of octasulfur features a triple point at approximately 115.2°C and 0.00026 MPa (26 Pa), where the monoclinic solid, liquid, and vapor phases coexist in equilibrium; this point marks the lower limit of the liquid-vapor boundary for the β-form. Due to the α-to-β transition at 95.5 °C, the rhombic (α) form does not directly coexist with liquid and vapor at this temperature.[27] The critical point occurs at 1041°C and 20.7 MPa, beyond which the distinction between liquid and vapor phases vanishes, defining the upper limit of the two-phase region. These parameters illustrate the energetic constraints on octasulfur's phase stability, with brief links to polymorph transitions such as rhombic to monoclinic observed in the solid region.[28]| Property | Value | Conditions | Source |
|---|---|---|---|
| Standard enthalpy of formation (ΔH_f°) | 0 kJ/mol | 298 K, solid | Standard Thermodynamic Values PDF |
| Standard Gibbs free energy of formation (ΔG_f°) | 0 kJ/mol | 298 K, solid | Standard Thermodynamic Values PDF |
| Standard entropy (S°) | 32.1 J/mol·K | 298 K, solid | Standard Thermodynamic Values PDF |
| Molar heat capacity (C_p) | 22.74 J/mol·K | 298 K, solid | Heat Capacity Table |
| Enthalpy of fusion (ΔH_fus) | 1.727 kJ/mol | 112.8°C | WebElements Sulfur Thermochemistry |
| Enthalpy of vaporization (ΔH_vap) | 45 kJ/mol | 444.6°C | WebElements Sulfur Thermochemistry |
| Triple point | 115.2°C, 0.00026 MPa | Monoclinic solid-liquid-vapor | Phase Diagram Explanation |
| Critical point | 1041°C, 20.7 MPa | Liquid-vapor | Periodic Table Data |
Production and Occurrence
Natural Sources
Octasulfur, the predominant allotrope of elemental sulfur, occurs naturally in various geological settings, primarily through volcanic and sedimentary processes. Native sulfur deposits are commonly associated with volcanic emissions, where sulfur vapors condense as elemental sulfur around fumaroles and vents, as observed in regions like the Gulf of Mexico salt domes and Yellowstone National Park. Hot springs also contribute to sulfur accumulation, as dissolved sulfur compounds precipitate upon cooling and degassing in hydrothermal systems. Sedimentary layers host significant reserves, often formed by the bacterial reduction of sulfate minerals followed by oxidation, with exogenic deposits in evaporitic sequences accounting for approximately 90% of the world's native sulfur reservoir. Historically, the Sicilian sulfur mines, part of the late Miocene Gypsum-Sulfur series, exemplified these sedimentary origins, yielding vast quantities from lagoonal evaporites in south and southwest Sicily until the early 20th century. In biological contexts, octasulfur forms through microbial activity in anoxic environments, particularly via the sulfur cycle involving sulfate-reducing bacteria (SRB) and sulfur-oxidizing bacteria. SRB, such as those in the Desulfobacterota phylum, reduce sulfate to hydrogen sulfide in oxygen-deprived sediments and waters, creating conditions for subsequent oxidation. At the oxic-anoxic interface, sulfur-oxidizing bacteria, including phototrophic species like Allochromatium vinosum, oxidize sulfide to produce octasulfur globules, often stored intracellularly or extracellularly as stable S8 aggregates. This process is prevalent in marine sediments, wetlands, and stratified lakes, where alternative electron acceptors like nitrate enable anoxic oxidation, contributing to natural sulfur deposition without requiring free oxygen. Octasulfur constitutes approximately 99% of naturally occurring elemental sulfur, existing as the stable orthorhombic α-S8 form under ambient conditions, with minor allotropes like monoclinic sulfur comprising the remainder. This dominance arises from the thermodynamic stability of the crown-shaped S8 ring, which minimizes strain energy compared to smaller rings or chains. Extraterrestrial occurrences of octasulfur further highlight its cosmic prevalence. On Jupiter's moon Io, volcanic activity ejects sulfur-rich plumes that condense into elemental sulfur deposits, rearranging into stable S8 rings on the cold surface, contributing to the moon's yellowish hues alongside SO2 frost. In comets, such as 67P/Churyumov-Gerasimenko, sulfur-bearing species including S8 have been detected in the coma through spectroscopic analysis by the Rosetta mission, likely originating from irradiated H2S ices in the nucleus. Recent analysis of samples from the asteroid Ryugu, returned by Japan's Hayabusa2 mission, confirmed octasulfur clusters via mass spectrometry, indicating primordial formation in aqueous-altered carbonaceous materials. Humans have utilized native sulfur since ancient times, with evidence of its extraction and application in Mesopotamia dating back to around 2500 BCE, where Sumerians employed it as a fumigant and pesticide for crops and insects. The molecular structure of octasulfur as an eight-membered ring was elucidated in the 19th century, with Hermann W. Vogel's 1895 determination of its molecular weight in solution providing key evidence for the S8 formula, building on earlier vapor density measurements.Industrial Production
The primary industrial method for producing octasulfur, the predominant allotrope of elemental sulfur, is the Claus process, which recovers sulfur from hydrogen sulfide (H₂S) contained in natural gas streams and refinery off-gases.[29] In this process, partial combustion of H₂S in air or oxygen produces sulfur dioxide (SO₂) and elemental sulfur in the first stage, followed by catalytic conversion where additional H₂S reacts with SO₂ to form more sulfur via the overall simplified reaction 2 H₂S + O₂ → 2 S + 2 H₂O.[30] The process typically achieves an overall sulfur recovery yield exceeding 95%, with modern configurations reaching 98-99% through multiple catalytic stages using alumina-based catalysts.[31] Historically, the Frasch process extracted native sulfur deposits from underground salt domes by injecting superheated water (around 160°C) to melt the sulfur, which was then pumped to the surface as a molten liquid.[32] Developed in the late 19th century, this method dominated U.S. production until the mid-20th century but has been largely discontinued in many regions post-2000 due to depletion of accessible deposits and the economic superiority of recovery processes from sour gases.[32] Today, Frasch mining persists only in limited operations, such as in Poland and Russia, contributing a minor fraction to global supply.[33] A significant portion of octasulfur originates as a byproduct from petroleum refining, where H₂S is removed from crude oil fractions via hydrodesulfurization, and from metal smelting operations processing sulfide ores of copper, zinc, and lead.[34] In petroleum refining, sulfur recovery is mandatory to meet environmental regulations, accounting for over 80% of global elemental sulfur output, while smelting contributes smaller but consistent volumes through similar H₂S-to-sulfur conversion.[34] These sources ensure a steady supply tied to fossil fuel and metallurgical industries. Following recovery, crude molten sulfur is refined to isolate high-purity α-octasulfur, the stable rhombic crystalline form, through methods such as vacuum distillation to remove volatile impurities or centrifugation to separate solids and water.[35] Distillation involves heating the melt to 200-300°C under reduced pressure, vaporizing sulfur while condensing it as pure α-crystals, whereas centrifugation uses high-speed separators to produce a clarified product with over 99.5% purity suitable for industrial use.[36] Global production of elemental sulfur, predominantly as octasulfur, reached approximately 85 million metric tons in 2023, with major contributors including China, the United States, and Russia.[37] Recent advancements in the Claus process since 2020 have focused on catalytic improvements, such as the adoption of alkali-promoted titania catalysts and sub-dewpoint operation, which enhance hydrolysis of sulfur compounds and boost overall efficiency to above 99.5% while reducing energy consumption.[38] These optimizations, including kinetic modeling for better reactor design, address challenges in treating low-H₂S acid gases and support higher recovery rates in integrated refinery operations.[39]Chemical Properties and Reactions
Reactivity Overview
Octasulfur (S₈), the predominant allotrope of elemental sulfur, exhibits high thermodynamic stability and chemical inertness at room temperature under ambient conditions, rendering it unreactive toward most reagents and poorly soluble in common solvents.[40] This stability arises from its symmetric crown-like structure with strong S–S bonds, allowing it to persist as a yellow solid without significant decomposition. However, reactivity markedly increases at elevated temperatures above approximately 100°C, where S₈ undergoes ring-opening and participates in processes such as homopolymerization into diradical chains, often requiring heating to 130–159°C for efficient transformation.[40] In S₈, each sulfur atom resides in the zero oxidation state (S⁰), but it readily shifts to higher states like +4 (as in SO₂) or +6 (as in sulfate species) during oxidative reactions, reflecting sulfur's versatile redox chemistry across oxidation states from -2 to +6.[40] The primary mode of S₈ activation involves ring-opening through cleavage of its S–S bonds, where the intact S₈ molecule functions as an electrophile susceptible to nucleophilic attack, particularly at the S² position.[40] Nucleophiles such as bases, sulfides, or amines initiate this process by breaking one or more bonds, generating reactive polysulfide intermediates (e.g., S₉²⁻ or S₅²⁻) that propagate further reactions.[40] This electrophilic behavior underpins S₈'s role in synthetic transformations, with mechanisms ranging from nucleophilic to radical pathways depending on conditions, though the molecule remains largely dormant without such activation.[41] In terms of redox properties, S₈ is readily reduced to polysulfides in aprotic or basic environments, forming stable radical anions like S₈⁻ as key intermediates during electrochemical processes.[42] Oxidation, conversely, converts S₈ to species such as SO₂, especially under combustive or oxidative conditions, highlighting its susceptibility to both cathodic and anodic shifts.[40] Solvent polarity significantly modulates this reactivity; polar media like DMSO, DMF, or water stabilize charged intermediates and polysulfides, enhancing solubility (e.g., up to 10⁵-fold in aqueous cyclodextrin solutions) and accelerating reactions such as thiol-mediated reductions compared to nonpolar solvents like toluene.[40][43]Specific Reactions
Octasulfur undergoes combustion in the presence of oxygen to form sulfur dioxide, a highly exothermic process central to sulfur's role in energy release during burning. The balanced reaction is given by
with a standard enthalpy change of approximately -297 kJ/mol per sulfur atom, totaling -2,376 kJ/mol for the octasulfur molecule.[44] This reaction proceeds vigorously at elevated temperatures, often exceeding 1,000°C, and is the primary pathway for sulfur oxidation in flames or industrial furnaces.[45]
Octasulfur reacts with alkali metals such as sodium at high temperatures to yield metal sulfides, demonstrating its reducing character toward electropositive elements. The reaction with sodium produces sodium sulfide via
typically requiring molten conditions or elevated temperatures around 300–500°C to overcome kinetic barriers and ensure complete ring opening.[46] This synthesis forms the basis for preparing binary sulfides used in further inorganic transformations, with the process being highly exothermic due to strong metal-sulfur bonding.[47]
Interaction of octasulfur with halogens like chlorine leads to halogenation products, highlighting its electrophilic behavior at sulfur atoms. With excess chlorine, the reaction forms sulfur dichloride according to
an exothermic process occurring at moderate temperatures (20–100°C) that proceeds via initial ring opening and subsequent chlorination.[48] This reaction is utilized in laboratory syntheses of sulfur halides, where the product SCl₂ serves as a chlorinating agent, though it decomposes readily if not handled under inert conditions.
Hydrogenation of octasulfur reduces it to hydrogen sulfide under catalytic conditions, illustrating its ability to act as a hydrogen acceptor in reductive environments. The transformation follows
facilitated by metal sulfide catalysts (e.g., MoS₂ or Ni-based) at temperatures of 300–400°C and elevated hydrogen pressures (typically 10–50 bar) to promote ring cleavage and stepwise reduction.[49] This reaction is relevant in hydroprocessing contexts, where sulfur removal from feeds occurs via similar mechanisms, though direct hydrogenation of elemental sulfur requires careful control to avoid side products like polysulfanes.[50]
Thermal decomposition of octasulfur above 500°C establishes an equilibrium favoring smaller sulfur species, driven by entropy gains from fragmentation. Smaller rings such as S₆ (cyclohexasulfur) and S₂ (disulfur) become prominent, with the reaction shifting toward products as temperature increases due to weakened S–S bonds. Studies reveal radical mechanisms underlying this process, where homolytic cleavage initiates chain propagation and recombination, influencing vapor-phase speciation in high-temperature environments.
In moist air, octasulfur undergoes slow, non-catalytic oxidation to sulfuric acid, reflecting its mild reactivity with atmospheric oxygen and water vapor over extended periods. This incomplete process initially forms sulfur dioxide, which further hydrolyzes and oxidizes to H₂SO₄, but proceeds at negligible rates without accelerators, often requiring years for significant conversion under ambient conditions.[51] The reaction highlights octasulfur's environmental persistence, as the uncatalyzed pathway yields trace acidic species that contribute to gradual material degradation.[52]