Question

Sulfur forms a wide variety of compounds in which it has +6 \(+4,+2,0,\) and -2 oxidation states. Give examples of sulfur \- compounds having each of these oxidation states.

Short answer

Examples of sulfur compounds with various oxidation states are as follows: Oxidation State +6: Sulfuric acid (H_2SO_4), sodium sulfate (Na_2SO_4), and benzenesulfonic acid (C_6H_5SO_3H). Oxidation State +4: Sulfur dioxide (SO_2), dimethyl sulfoxide ((CH_3)_2SO), and thiosulfate ion (S_2O_3^{-2}). Oxidation State +2: Hydrogen sulfide (H_2S), iron sulfide (FeS), and dimethyl sulfide ((CH_3)_2S). Oxidation State 0: Elemental sulfur (S_8), disulfur (S_2), and thiophene (C_4H_4S). Oxidation State -2: Sodium persulfide (Na_2S_2) and diphenyl disulfide ((C_6H_5S)_2).

Step-by-step solution

Definition of Oxidation States

The oxidation state is the hypothetical charge that an atom would have if all its bonds to elements of different atoms were 100% ionic in nature. It is a way of keeping track of electrons and simplifying redox reactions. Oxidation states can range from positive to negative values, including zero.

Sulfur Compounds with Oxidation State +6

Compounds in which sulfur exhibits an oxidation state of +6 include inorganic compounds like sulfuric acid (H_2SO_4), sulfates (e.g., sodium sulfate, Na_2SO_4), and some organic compounds containing sulfonyl groups (e.g., benzenesulfonic acid, C_6H_5SO_3H). In these compounds, sulfur is covalently bonded to more electronegative atoms like oxygen and loses six electrons.

Sulfur Compounds with Oxidation State +4

In compounds where sulfur has an oxidation state of +4, sulfur loses four electrons and forms covalent bonds with either oxygen or other atoms in the compound. Examples of such compounds include sulfur dioxide (SO_2), sulfoxides in organic chemistry (e.g., dimethyl sulfoxide, (CH_3)_2SO), and thiosulfate ion (S_2O_3^{-2}).

Sulfur Compounds with Oxidation State +2

Compounds where sulfur exhibits an oxidation state of +2 include hydrogen sulfide (H_2S), sulfides (e.g., iron sulfide, FeS), and organic thioethers (e.g., dimethyl sulfide, (CH_3)_2S). In these compounds, sulfur loses two electrons to become the positively charged part of the compound.

Sulfur Compounds with Oxidation State 0

When sulfur exhibits an oxidation state of 0, it means that it is in its elemental form and has not lost or gained any electrons. Examples include elemental sulfur (S_8) and sulfur-containing molecules like disulfur (S_2) or sulfur ring structures (e.g., thiophene, C_4H_4S).

Sulfur Compounds with Oxidation State -2

When sulfur has an oxidation state of -2, it means that it has gained two electrons in the compound. Compounds in which sulfur exhibits an oxidation state of -2 include metal persulfides (e.g., sodium persulfide, Na_2S_2) and organic compounds containing persulfide groups (e.g., diphenyl disulfide, (C_6H_5S)_2).

Key concepts

Sulfuric Acid

Sulfuric acid (H₂SO₄) is a paramount example of sulfur in its +6 oxidation state, symbolizing a highly oxidized form. This colorless, viscous liquid is known for its strong acidic properties and ability to cause severe burns. As a strong acid, it easily donates protons (H⁺) in reactions. It plays a crucial role in synthesizing a wide range of chemicals, and in processes like battery acid, and fertilizer production.

Sulfuric acid is made up of hydrogen, sulfur, and oxygen, where sulfur is centrally bonded to four oxygen atoms. Two of these are double-bonded, known as oxo groups, and the other two are single-bonded through hydroxide groups. The secure bonding with the highly electronegative oxygen atoms ensures that sulfur maintains a +6 oxidation state, allowing it to act as an oxidizing agent in many reactions.

Sulfides

Sulfides are compounds that showcase sulfur in a -2 oxidation state, typically combined with metals to form an array of minerals and materials. A common example would be iron sulfide (FeS), which occurs naturally and can be synthesized for various uses, from pigments to semiconductor technology.

In sulfides, sulfur atoms possess two additional electrons, which is why their oxidation state is negative. This anionic form of sulfur (S^2-) is fundamentally different from sulfuric acid's highly oxidized version. Sulfides exhibit different physical and chemical properties, such as non-polar characteristic and lesser solubility in water, compared to sulfuric acid. Understanding the formation and reactivity of sulfides is integral in fields like geology, materials science, and industrial chemistry.

Sulfur Compounds

Sulfur compounds possess a diverse array of oxidation states, thus creating a broad spectrum of chemical behavior and applications. The ubiquitous nature of sulfur compounds is evident from their presence in fuels, drugs, and even in the biological systems as essential amino acids. The oxidation state of sulfur in these compounds is a key determinant of their chemical properties, whether they are used as reducing agents, oxidizing agents, or as part of more complex biochemical pathways.

The ability of sulfur to form stable compounds with both metals and nonmetals, including hydrogen, oxygen, and other chalcogens, is facilitated by its position in the periodic table. This affords sulfur significant versatility, contributing to the formation of compounds like sulfates, sulfides, thiosulfates, and polysulfides, each demonstrating unique reactivity and uses. Delving into the chemistry of sulfur compounds enhances our understanding of environmental processes and offers insights into developing new materials and pharmaceuticals.

Redox Reactions

Redox reactions are fundamental chemical processes where electrons are transferred between substances, consequently changing their oxidation states. These reactions are integral for understanding how sulfur compounds operate and react in various environments. Oxidation involves the loss of electrons, while reduction entails the gain of electrons. In the context of sulfur, we witness its versatility through redox reactions that include elemental sulfur being oxidized to form sulfates or sulfides being oxidized to sulfuric acid.

Exploring redox reactions aids in grasping the underlying principles governing battery function, metallurgical processes, and biochemical energy generation within living organisms. For example, the use of sulfides such as FeS2 (pyrite) in batteries relies on redox reactions to release energy. Similarly, within biological systems, redox reactions involving sulfur compounds are pivotal for cellular respiration and maintaining life. Brightening the concepts of oxidation states and redox reactions serves as a foundation for future studies in chemistry and related disciplines.