Sulfur

Properties of Sulfur

Sulfur burns with blue flame, is insoluble in water, and forms polyatomic allotropes.

Learning Objectives

Describe the properties of sulfur.

Key Takeaways

Key Points

  • Sulfur forms polyatomic molecules with different chemical formulas. Its best-known allotrope is octasulfur, S8, which is a soft, yellow solid with faint odor. Due to changes in intermolecular interactions, it undergoes phase changes from α-octasulfur to β-polymorph to γ-sulfur at high temperatures.
  • At temperatures higher than boiling point for octasulfur, depolymerization occurs. Molten sulfur has a dark red color above 200 °C. Different allotropes have different densities of about 2 g/cm3. Stable allotropes are excellent electrical insulators.
  • Sulfur burns with a blue flame, forming sulfur dioxide with suffocating odor. It is insoluble in water but soluble in carbon disulfide. S+4, S6+ are more common than S2+. Higher ionization states exist only with strong oxidants such as fluorine, oxygen, and chlorine.

Key Terms

  • depolymerization: The decomposition of a polymer into smaller fragments.
  • octasulfur: The most common allotrope of sulfur (S8) containing eight atoms in a ring.
  • allotrope: Any form of a pure element that has a distinctly different molecular structure.

Octasulfur

Sulfur is found is different polyatomic allotropic forms. The best-known allotrope is octasulfur, cyclo-S8. Octasulfur is a soft, bright-yellow solid with only a faint odor, similar to that of matches. It melts at 115.21 °C, boils at 444.6 °C, and sublimes easily.

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Cyclooctasulfur: The structure of the cyclooctasulfur molecule, S8.

At 95.2 °C, below its melting temperature, cyclooctasulfur changes from α-octasulfur to the β-polymorph. The structure of the S8 ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotropic form again, turning from β-octasulfur to γ-sulfur. Again, this is accompanied by a lower density but increased viscosity due to the formation of polymers. At even higher temperatures, however, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C. The density of sulfur is about 2 g/cm3, depending on the allotrope. All of sulfur’s stable allotropes are excellent electrical insulators.

Chemical Properties of Sulfer

Sulfur burns with a blue flame, concomitant with formation of sulfur dioxide, notable for its peculiar suffocating odor. Sulfur is insoluble in water but soluble in carbon disulfide and, to a lesser extent, in other nonpolar organic solvents, such as benzene and toluene. The first and the second ionization energies of sulfur are 999.6 and 2252 kJ/mol, respectively. Despite such figures, S2+ is rare, with S+4 and S6+ being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ/mol. The magnitude of the figures is caused by electron transfer between orbitals; these states are only stable with strong oxidants such as fluorine, oxygen, and chlorine.

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Melting and burning sulfur: Sulfur burns with blue flames and forms blood-red liquid when it melts.

Sulfur Compounds

Sulfur forms stable compounds with most elements except the noble gases.

Learning Objectives

Discuss several examples of sulfur compounds.

Key Takeaways

Key Points

  • Hydrogen sulfide is mildly acidic in water and is extremely toxic. Sulfur can form chains with itself (catenation). Polysulfides are formed by reduction of elemental sulfur. Polysulfanes are protonated polysulfides. Reduction of sulfur gives sulfide salts.
  • Burning sulfur forms the principal sulfur oxides. The sulfur oxides form numerous oxyanions, which are related to numerous acids. Oleum is a solution of pyrosulfuric acid and sulfuric acid. Peroxides convert sulfur into unstable sulfoxides.
  • Sulfur compounds with halogens include sulfur hexafluoride, sulfur dichloride, and chlorosulfuric acid. Thionyl chloride is a common reagent in organic synthesis. Tetrasulfur tetranitride and thiocyanates are compounds of sulfur and nitrogen. Phosphorus sulfides are numerous.
  • The principal ores of many metals are sulfides. They are formed by the reaction of hydrogen sulfide with metal salts. Tarnishing is the process of metal corrosion by sulfur.
  • Sulfur-containing organic compounds include thiols (the sulfur analogs of alcohols) and thioethers (the sulfur analogs of ethers ). Compounds with carbon–sulfur bonds are uncommon. Organosulfur compounds are responsible for the some of the unpleasant odors of decaying organic matter.
  • Sulfur-sulfur bonds are a structural component of proteins, providing them with rigidity. Vulcanization is the process of heating rubber and sulfur until disulfide bridges form between isoprene units of the polymer. This increases rigidity of rubber.

Key Terms

  • catenation: The ability of a few elements, most especially carbon, to yield chains and rings by forming covalent bonds with atoms of the same element.
  • vulcanization: A process by which rubber is hardened using heat and sulfur.

Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases. For some organic sulfur compounds, smell depends on their concentration. The sulfur-containing monoterpenoid grapefruit mercaptan has the characteristic scent of grapefruit in small concentrations, but has a unpleasant thiol odor at larger concentrations.

Reactions with Hydrogen

Treatment of sulfur with hydrogen produces hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:

[latex]\text{H}_2\text{S} \leftrightarrow \text{HS}^- + \text{H}^+[/latex]

Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, because they inhibit the oxygen -carrying capacity of hemoglobin and certain cytochromes in a manner similar to cyanide and azide.

Reduction of Sulfur

Reduction of elemental sulfur produces polysulfides, which consist of chains of sulfur atoms terminated with S centers:

[latex]2 \text{Na} + \text{S}_8 \rightarrow \text{Na}_2\text{S}_8[/latex]

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Lapis lazuli: Lapis lazuli owes its blue color to a sulfur radical.

This reaction highlights arguably the single most distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions gives the polysulfanes, H2Sx where x = 2, 3, and 4.

Ultimately, reduction of sulfur gives sulfide salts:

[latex]16 \text{Na} + \text{S}_8 \rightarrow 8 \text{Na}_2\text{S}[/latex]

The interconversion of these species is used in sodium-sulfur batteries.

The radical anion S3 gives the blue color of the mineral lapis lazuli. With very strong oxidants, S8 can be oxidized, for example, to give bicyclic S82+.

Sulfur Oxides

The principal sulfur oxides are obtained by burning sulfur:

[latex]\text{S} + \text{O}_2 \rightarrow \text{SO}_2[/latex],

[latex]2 \text{SO}_2 + \text{O}_2 \rightarrow 2 \text{SO}_3[/latex]

Other oxides are known—sulfur monoxide and disulfur mono- and dioxides—but they are unstable. The sulfur oxides form numerous oxyanions with the formula SOn2–. Sulfur dioxide and sulfites (SO32−) are related to the unstable sulfurous acid (H2SO3). Sulfur trioxide and sulfates (SO42−) are related to sulfuric acid. Sulfuric acid and SO3 combine to give oleum, a solution of pyrosulfuric acid (H2S2O7) in sulfuric acid.

Peroxides convert sulfur into unstable compounds such as S8O, a sulfoxide. Peroxymonosulfuric acid (H2SO5) and peroxydisulfuric acids (H2S2O8) are made from the action of SO3 on concentrated H2O2, and H2SO4 on concentrated H2O2, respectively. Thiosulfate salts(S2O32−), sometimes referred as “hyposulfites,” are used in photographic fixing (HYPO) and as reducing agents. These salts feature sulfur in two oxidation states. Sodium dithionite, (S2O42−) contains the more highly reducing dithionite anion; sodium dithionate (Na2S2O6) is the first member of the polythionic acids (H2SnO6), where n can range from 3 to many.

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Peroxydisulfuric acid: Peroxydisulfuric acid (H2S2O8) is made from the action of H2SO4 on concentrated H2O2.

Sulfur Halides

There are two main sulfur fluorides. Sulfur hexafluoride is a dense gas used as a nonreactive and nontoxic propellant. Sulfur tetrafluoride is a rarely used organic reagent that is highly toxic. Their chlorinated analogs are sulfur dichloride and sulfur monochloride. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl2) is a common reagent in organic synthesis.

Sulfur-Nitrogen Compounds

An important S–N compound is the cage tetrasulfur tetranitride (S4N4). Heating this compound gives polymeric sulfur nitride ((SN)x), which has metallic properties even though it does not contain any metal atoms. Thiocyanates contain the SCN group. Oxidation of thiocyanate gives thiocyanogen, (SCN)2 with the connectivity NCS–SCN. Phosphorus sulfides are numerous.

Sulfur with Metals

The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides. These materials tend to be dark-colored semiconductors that are not readily attacked by water or even many acids. They are formed by the reaction of hydrogen sulfide with metal salts. The mineral galena (PbS) was the first demonstrated semiconductor. It was used as a signal rectifier in the cat’s whiskers of early crystal radios. Upgrading these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals via the process called tarnishing.

Organic Compounds with Sulfur

The following are some of the main classes of sulfur-containing organic compounds:

  • Thiols or mercaptans (as they are mercury capturers as chelators) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
  • Thioethers are the sulfur analogs of ethers. Sulfonium ions have three groups attached to a cationic sulfur center. Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine organic sulfur cycle.
  • Sulfoxides and sulfones are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, dimethyl sulfoxide, is a common solvent; a common sulfone is sulfolane. Sulfonic acids are used in many detergents.
  • Compounds with carbon–sulfur bonds are uncommon except for carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. Unlike carbon monoxide, carbon monosulfide is only stable as a dilute gas, as in the interstellar medium. Organosulfur compounds are responsible for the some of the unpleasant odors of decaying organic matter.

Vulcanization

Sulfur-sulfur bonds are a structural component to stiffen rubber, similar to the biological role of disulfide bridges in rigidifying proteins. In the most common type of industrial “curing” or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber until chemical reactions form disulfide bridges between isoprene units of the polymer. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.