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Chapter 20: Problem 85

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

Expert verified
Examples of sulfur compounds with various oxidation states are: 1. For an oxidation state of +6, sulfuric acid (H₂SO₄) is an example. 2. For an oxidation state of +4, sulfur dioxide (SO₂) is an example. 3. For an oxidation state of +2, hydrogen sulfide (H₂S) is an example. 4. For an oxidation state of 0, elemental sulfur (S) is an example. 5. For an oxidation state of -2, a sulfide ion (S²⁻) found in compounds like iron sulfide (FeS) is an example.

Step by step solution

01

Oxidation state: +6

One example of a sulfur compound with an oxidation state of +6 is sulfuric acid (H₂SO₄). In this compound, sulfur is bonded to four oxygen atoms and has an overall charge of +6.
02

Oxidation state: +4

An example of a sulfur compound with an oxidation state of +4 is sulfur dioxide (SO₂). In this compound, sulfur is bonded to two oxygen atoms and has an overall charge of +4.
03

Oxidation state: +2

One example of a sulfur compound with an oxidation state of +2 is hydrogen sulfide (H₂S). In this compound, sulfur is bonded to two hydrogen atoms and has an overall charge of +2.
04

Oxidation state: 0

An example of a sulfur compound with an oxidation state of 0 is elemental sulfur (S). In its elemental form, sulfur is not bonded to any other atoms and thus has an oxidation state of 0.
05

Oxidation state: -2

One example of a sulfur compound with an oxidation state of -2 is a sulfide ion (S²⁻), which can be found in compounds like iron sulfide (FeS). In this case, sulfur has an overall charge of -2 as it gains two electrons from the metal atom it is bonded to.

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Most popular questions from this chapter

Many structures of phosphorus-containing compounds are drawn with some \(\mathrm{P}=\mathrm{O}\) bonds. These bonds are not the typical \(\pi\) bonds we've considered, which involve the overlap of two \(p\) orbitals. Instead, they result from the overlap of a \(d\) orbital on the phosphorus atom with a \(p\) orbital on oxygen. This type of \(\pi\) bonding is sometimes used as an explanation for why \(\mathrm{H}_{3} \mathrm{PO}_{3}\) has the first structure below rather than the second: Draw a picture showing how a \(d\) orbital and a \(p\) orbital overlap to form a \(\pi\) bond.

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Explain why \(\mathrm{HF}\) is a weak acid, whereas \(\mathrm{HCl}, \mathrm{HBr}\), and \(\mathrm{HI}\) are all strong acids.

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