Question

Arrange the following species in order of increasing oxidation number of the sulfur atom: (a) \(\mathrm{H}_{2} \mathrm{~S},(\mathrm{~b}) \mathrm{S}_{8}\) (c) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) \(\mathrm{S}^{2-}\) (e) HS \(^{-}\), (f) \(\mathrm{SO}_{2},(\mathrm{~g}) \mathrm{SO}_{3}\)

Short answer

Order: \( \mathrm{S}^{2-} = \mathrm{HS}^{-} = \mathrm{H}_{2}\mathrm{S} < \mathrm{S}_{8} < \mathrm{SO}_{2} < \mathrm{H}_{2}\mathrm{SO}_4 = \mathrm{SO}_3 \).

Step-by-step solution

Understand Oxidation Number Concept

Oxidation numbers are a way to keep track of electrons in atoms, especially during chemical reactions. It is a hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic.

Identify oxidation numbers in individual species

Calculate or determine the oxidation number of sulfur in each species presented: (a) In \( \mathrm{H}_{2}\mathrm{S} \): Hydrogen is \(+1\) and there are two hydrogens, thus sulfur is \(-2\).(b) In \( \mathrm{S}_{8} \): This is elemental sulfur, so its oxidation number is \(0\).(c) In \( \mathrm{H}_{2}\mathrm{SO}_{4} \): Hydrogen is \(+1\) (total for both hydrogens is \(+2\)), oxygen is \(-2\) (total for four oxygens is \(-8\)), thus sulfur is \(+6\).(d) In \( \mathrm{S}^{2-} \): The oxidation number of sulfur is equivalent to its charge, which is \(-2\).(e) In HS \(^{-}\): Hydrogen is \(+1\), so sulfur must be \(-2\) to balance the net charge of \(-1\).(f) In \( \mathrm{SO}_{2} \): Oxygen is \(-2\) and there are two oxygens (total \(-4\)), so sulfur is \(+4\).(g) In \( \mathrm{SO}_{3} \): Oxygen is \(-2\) and there are three oxygens (total \(-6\)), so sulfur is \(+6\).

Arrange species by oxidation number

List the determined oxidation numbers of sulfur for comparison:- (d) \( \mathrm{S}^{2-}: -2 \)- (e) \( \mathrm{HS}^{-}: -2 \)- (a) \( \mathrm{H}_2\mathrm{S}: -2 \)- (b) \( \mathrm{S}_8: 0 \)- (f) \( \mathrm{SO}_2: +4 \)- (c) \( \mathrm{H}_2\mathrm{SO}_4: +6 \)- (g) \( \mathrm{SO}_3: +6 \)Now, arrange them in increasing order: \( \mathrm{S}^{2-} = \mathrm{HS}^{-} = \mathrm{H}_{2}\mathrm{S} < \mathrm{S}_{8} < \mathrm{SO}_{2} < \mathrm{H}_{2}\mathrm{SO}_4 = \mathrm{SO}_3 \).

Key concepts

Sulfur Compounds

Sulfur is a versatile element found in various compounds, playing significant roles in biology, industry, and the environment. Sulfur compounds can range from simple molecules, like hydrogen sulfide (\( \mathrm{H}_2\mathrm{S} \)), to complex acids, such as sulfuric acid (\( \mathrm{H}_2\mathrm{SO}_4 \)). The wide variety of sulfur compounds is due to sulfur's ability to form multiple oxidation states, allowing it to bond with many different elements.

Some key properties of sulfur compounds include:

• Hydrogen sulfide (\( \mathrm{H}_2\mathrm{S} \)): A simple compound where sulfur is bonded to hydrogen. It has a distinct rotten egg smell and is often associated with natural gas and volcanic activities.
• Sulfur dioxide (\( \mathrm{SO}_2 \)) and sulfur trioxide (\( \mathrm{SO}_3 \)): Important compounds used in the production of sulfuric acid. They are also byproducts of combustion in fossil fuels and can contribute to air pollution.
• Sulfide ions (\( \mathrm{S}^{2-} \)): These are often found in metal ores and are essential in the mining industry for extracting metals.
• Elemental sulfur (\( \mathrm{S}_8 \)): Exists in various allotropes, with \( \mathrm{S}_8 \) being the most stable and common form. It is a yellow solid used in fertilizers, insecticides, and vulcanization of rubber.

Understanding these key sulfur compounds and their interactions in nature and industry is crucial, as it helps us manage essential processes in environmental and chemical engineering strategies.

Oxidation States

An oxidation state, or oxidation number, is an integer that represents the electron count balance around an atom within a compound. It helps chemists understand how electrons are transferred in chemical reactions. It's important to familiarize oneself with how these numbers are assigned, especially with elements like sulfur that can exhibit different oxidation states.

For sulfur compounds listed in the exercise:

• Sulfide ion (\( \mathrm{S}^{2-} \)): Has an oxidation state of \,\(-2\), meaning it has gained two electrons compared to its elemental form.
• Hydrogen sulfide (\( \mathrm{H}_2\mathrm{S} \)) and HS^-: Similarly, these have sulfur oxidation states of \,\(-2\).
• Elemental sulfur (\( \mathrm{S}_8 \)): With an oxidation state of \,\(0\), as it is in its native form, not bonded with any electronegative elements.
• Sulfur dioxide (\( \mathrm{SO}_2 \)): Here, sulfur has an oxidation state of \,\(+4\).
• Sulfur trioxide (\( \mathrm{SO}_3 \)) and sulfuric acid (\( \mathrm{H}_2\mathrm{SO}_4 \)): Both present sulfur in its highest common oxidation state of \,\(+6\).

The concept of oxidation states is vital in balances during chemical reactions, allowing prediction of molecule interaction behavior.

Chemical Reactions

Chemical reactions are processes where reactants transform into products, often involving changes in oxidation states and the rearrangement of atoms. Key to these reactions is the transfer of electrons, which is tracked through oxidation numbers. This tracking helps chemists to predict and control reactions.

When considering sulfur in chemical reactions, note:

• Redox reactions: Reactions where oxidation and reduction occur simultaneously. Sulfur compounds frequently undergo such processes, as sulfur can easily change its oxidation state. For instance, sulfur dioxide oxidized to sulfur trioxide by gaining oxygen atoms.
• Acid-base reactions: Sulfuric acid (\( \mathrm{H}_2\mathrm{SO}_4 \)) is a strong acid and participates actively in donating protons in reactions, showcasing its utility in industry for processes like lead-acid batteries and fertilizer production.
• Combustion reactions: Involving sulfur compounds release energy, for instance, burning coal releases sulfur dioxide, a major air pollutant that can lead to the formation of acid rain.

Understanding sulfur's role in these reactions is pivotal in fields such as environmental science, industrial chemistry, and pollution control. By controlling sulfur's oxidation states and reaction conditions, we can optimize these chemical reactions for beneficial applications.