22.6 The Other Group 6A Elements: S, Se, Te, and Po

In addition to oxygen, the other group 6A elements are sulfur, selenium, tellurium, and polonium. In this section we will survey the properties of the group as a whole and then examine the chemistry of sulfur, selenium, and tellurium. We will not say much about polonium, which has no stable isotopes and is found only in minute quantities in radium-containing minerals.

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General Characteristics of the Group 6A Elements

The group 6A elements possess the general outer-electron configuration ns2np4, where n has values ranging from 2 through 6. Thus, these elements may attain a noble-gas electron configuration by the addition of two electrons, which results in a -2 oxidation state. Because the group 6A elements are nonmetals, this is a common oxidation state. Except for oxygen, however, the group 6A elements are also commonly found in positive oxidation states up to +6, and they can have expanded valence shells. Thus, compounds such as SF6, SeF6, and TeF6 occur in which the central atom is in the +6 oxidation state with more than an octet of valence electrons.

Table 22.4 summarizes some of the more important properties of the atoms of the group 6A elements. In most of the properties listed in Table 22.4, we see a regular variation as a function of atomic number. For example, atomic and ionic radii increase and ionization energies decrease, as expected, as we move down the family.

Occurrences and Preparation of S, Se, and Te

Large underground deposits are the principal source of elemental sulfur. The Frasch process, illustrated in Figure 22.23, is used to obtain the element from these deposits. The method is based on the low melting point and low density of sulfur. Superheated water is forced into the deposit, where it melts the sulfur. Compressed air then forces the molten sulfur up a pipe to the surface above, where the sulfur cools and solidifies.

Figure 22.23 Mining of sulfur by the Frasch process. The process is named after Herman Frasch, who invented it in the early 1890s. The process is particularly useful for recovering sulfur from deposits located under quicksand or water.

Sulfur also occurs widely as sulfide and sulfate minerals. Its presence as a minor component of coal and petroleum poses a major problem. Combustion of these "unclean" fuels leads to serious sulfur oxide pollution. Much effort has been directed at removing this sulfur, and these efforts have increased the availability of sulfur. The sale of this sulfur helps partially to offset the costs of the desulfurizing processes and equipment. About half the sulfur used in the United States each year is produced by means other than the Frasch process.

Selenium and tellurium occur in rare minerals such as Cu2Se, PbSe, Ag2Se, Cu2Te, PbTe, Ag2Te, and Au2Te. They also occur as minor constituents in sulfide ores of copper, iron, nickel, and lead.

Properties and Uses of Sulfur, Selenium, and Tellurium

As we normally encounter it, sulfur is yellow, tasteless, and nearly odorless. It is insoluble in water and exists in several allotropic forms. The thermodynamically stable form at room temperature is rhombic sulfur, which consists of puckered S8 rings, as shown in Figure 22.24. When heated above its melting point (113°C), sulfur undergoes a variety of changes. The molten sulfur first contains S8 molecules and is fluid because the rings readily slip over one another. Further heating of this straw-colored liquid causes the rings to break; the fragments then join to form very long molecules that can become entangled. The sulfur consequently becomes highly viscous. This change is marked by a color change to dark reddish brown (Figure 22.25). Further heating breaks the chains, and the viscosity again decreases.

Figure 22.24 The common yellow crystalline form of rhombic sulfur consists of S8 molecules. These molecules are puckered rings of S atoms.

Most of the 1.3 1010 kg (14 million tons) of sulfur produced in the United States each year is used in the manufacture of sulfuric acid. Sulfur is also used in vulcanizing rubber, a process that toughens rubber by introducing cross-linking between polymer chains.

The most stable allotropes of both selenium and tellurium are crystalline substances containing helical chains of atoms, as illustrated in Figure 22.26. Each atom of the chain is close to atoms in adjacent chains, and it appears that some sharing of electron pairs between these atoms occurs.

Figure 22.26 Portion of the structure of crystalline selenium. The dashed lines represent weak bonding interactions between atoms in adjacent chains. Tellurium has the same structure.

The electrical conductivity of selenium is very low in the dark, but increases greatly upon exposure to light. This property of the element is utilized in photoelectric cells and light meters. Photocopiers also depend on the photoconductivity of selenium. Photocopy machines contain a belt or drum coated with a film of selenium. This drum is electrostatically charged and then exposed to light reflected from the image being photocopied. The electric charge drains from the selenium where the selenium is made conductive by exposure to light. A black powder (the "toner") sticks only to the areas that remain charged. The photocopy is made when the toner is transferred to a sheet of plain paper, which is heated to fuse the toner to the paper.

Sulfides

Sulfur forms compounds by direct combination with many elements. When the element is less electronegative than sulfur, sulfides, which contain S2–, form. For example, iron(II) sulfide, FeS, forms by direct combination of iron and sulfur. Many metallic elements are found in the form of sulfide ores, for example, PbS (galena) and HgS (cinnabar). A series of related ores containing the disulfide ion, S22– (analogous to the peroxide ion), are known as pyrites. Iron pyrite, FeS2, occurs as golden yellow cubic crystals (Figure 22.27). Because it has been occasionally mistaken for gold by miners, it is often called "fool's gold."

One of the most important sulfides is hydrogen sulfide, H2S. This substance is not normally produced by direct union of the elements because it is unstable at elevated temperature and decomposes into the elements. It is normally prepared by action of dilute acid on iron(II) sulfide:

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Hydrogen sulfide is often used in the laboratory for qualitative analysis of certain metal ions.

One of hydrogen sulfide's most readily recognized properties is its odor; H2S is largely responsible for the offensive odor of rotten eggs. Hydrogen sulfide is actually quite toxic. Fortunately, our noses are able to detect H2S in extremely low, nontoxic concentrations. Sulfur-containing organic molecules, which are similarly odoriferous, are added to natural gas to give it a detectable odor.

Oxides, Oxyacids, and Oxyanions of Sulfur

Sulfur dioxide is formed when sulfur is combusted in air; it has a choking odor and is poisonous. The gas is particularly toxic to lower organisms, such as fungi, and is consequently used for sterilizing dried fruit and wine. At 1 atm pressure and room temperature SO2 dissolves in water to produce a solution of about 1.6 M concentration. The SO2(aq) solution is acidic, and we describe it as sulfurous acid, H2SO3(aq). Sulfurous acid is a diprotic acid:

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Salts of SO32– (sulfites) and HSO3 (hydrogen sulfites or bisulfites) are well known. Small quantities of Na2SO3 or NaHSO3 are used as food additives to prevent bacterial spoilage. Because some people are extremely allergic to sulfites, all food products with sulfites must now carry a warning disclosing their presence.

Although combustion of sulfur in air produces mainly SO2, small amounts of SO3 are also formed. The reaction produces mainly SO2 because the activation-energy barrier for further oxidation to SO3 is very high unless the reaction is catalyzed. Sulfur trioxide is of great commercial importance because it is the anhydride of sulfuric acid. In the manufacture of sulfuric acid, SO2 is first obtained by burning sulfur. The SO2 is then oxidized to SO3, using a catalyst such as V2O5 or platinum. The SO3 is dissolved in H2SO4 because it does not dissolve quickly in water (Equation 22.49). The H2S2O7 formed in this reaction, called pyrosulfuric acid, is then added to water to form H2SO4, as shown in Equation 22.50:

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Commercial sulfuric acid is 98 percent H2SO4. It is a dense, colorless, oily liquid that boils at 340°C. Sulfuric acid has many useful properties: It is a strong acid, a good dehydrating agent, and a moderately good oxidizing agent. Its dehydrating ability is demonstrated in Figure 22.28.

Year after year, the production of sulfuric acid is the largest of any chemical produced in the United States. About 4.3 1010 kg (48 million tons) are produced annually in this country. Sulfuric acid is employed in some way in almost all manufacturing. Consequently, its consumption is considered a standard measure of industrial activity.

Sulfuric acid is classified as a strong acid. However, only the first hydrogen in sulfuric acid is completely ionized in aqueous solution. The second hydrogen ionizes only partially:

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Consequently, sulfuric acid forms two series of compounds: sulfates and bisulfates (or hydrogen sulfates). Bisulfate salts are common components of the "dry acids" used for adjusting the pH of swimming pools and hot tubs; they are also components of many toilet bowl cleaners.

Related to the sulfate ion is the thiosulfate ion, S2O32–, formed by boiling an alkaline solution of SO32– with elemental sulfur:

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The term thio indicates substitution of sulfur for oxygen. The structures of the sulfate and thiosulfate ions are compared in Figure 22.29. When acidified, the thiosulfate ion decomposes to form sulfur and H2SO3.

Figure 22.29 Comparison of the structures of (a) the sulfate, SO42–, and (b) the thiosulfate, S2O32–, ions.

The pentahydrated salt of sodium thiosulfate, Na2S2O3 5H2O, known as "hypo," is used in photography. Photographic film consists of a suspension of microcrystals of AgBr in gelatin. When exposed to light, some AgBr decomposes, forming very small grains of silver. When the film is treated with a mild reducing agent (the "developer"), Ag+ ions in AgBr near the silver grains are reduced, forming an image of black metallic silver. The film is then treated with sodium thiosulfate solution to remove the unexposed AgBr. The thiosulfate ion reacts with AgBr to form a soluble silver thiosulfate complex:

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This step in the process is called "fixing." Thiosulfate ion is also used in quantitative analysis as a reducing agent for iodine:

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