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Chapter 22: Problem 88

Explain why \(\mathrm{SO}_{2}\) can be used as a reducing agent but \(\mathrm{SO}_{3}\) cannot.

Short Answer

Expert verified
SO2 can be used as a reducing agent because it has an oxidation state of +4 and a bent molecular structure with a lone pair of electrons on the sulfur atom, making it more reactive and able to donate electrons. In contrast, SO3 has an oxidation state of +6 (its highest possible state), a trigonal planar structure with no lone pairs on the sulfur atom, making it less reactive and unable to donate electrons. Therefore, SO3 cannot be used as a reducing agent.

Step by step solution

01

Understanding redox reactions and reducing agents

A reducing agent is a substance that donates electrons to another substance in a redox (reduction-oxidation) reaction, and in the process, gets oxidized itself. In other words, reducing agents make it possible for another substance to be reduced (i.e., to gain electrons). Therefore, a good reducing agent should have electrons that can be easily removed.
02

Comparing oxidation state of sulfur in SO2 and SO3

In SO2 (sulfur dioxide), sulfur has an oxidation state of +4, whereas, in SO3 (sulfur trioxide), sulfur is in its highest possible oxidation state of +6. A substance that is already in its highest oxidation state cannot be oxidized further and therefore cannot act as a reducing agent.
03

Understanding key differences between SO2 and SO3

Apart from the difference in sulfur's oxidation state, the two compounds also differ in structure and stability. SO2 is a bent molecule with a lone pair of electrons on the sulfur atom, which makes it more reactive and able to donate electrons. In contrast, SO3 has a trigonal planar structure with no lone pairs on the sulfur atom, which makes it less reactive and unable to donate electrons.
04

Explaining why SO2 can be used as a reducing agent

Given its electron configuration and oxidation state, SO2 can donate electrons to other molecules. For example, SO2 can reduce oxygen (O2) to form SO3: \(2\mathrm{SO}_2 + \mathrm{O}_2 \rightarrow 2\mathrm{SO}_3\) In this reaction, SO2 gets oxidized to SO3 (sulfur's oxidation state increases from +4 to +6), while oxygen gets reduced (gains electrons). Since SO2 can be oxidized and donate electrons, it can act as a reducing agent.
05

Explaining why SO3 cannot be used as a reducing agent

On the other hand, SO3 is already in its highest possible oxidation state of +6, which means it cannot donate electrons or be oxidized further. Additionally, its structure, with no lone pairs on the sulfur atom, makes it less reactive and unable to participate in redox reactions as a reducing agent. In conclusion, SO2 can be used as a reducing agent because it is able to donate electrons and be oxidized, while SO3 cannot due to its highest oxidation state and less reactive structure.

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Most popular questions from this chapter

Both dimethylhydrazine, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NNH}_{2}\), and methylhydrazine, \(\mathrm{CH}_{3} \mathrm{NHNH}_{2}\), have been used as rocket fuels. When dinitrogen tetroxide \(\left(\mathrm{N}_{2} \mathrm{O}_{4}\right)\) is used as the oxidizer, the products are \(\mathrm{H}_{2} \mathrm{O}, \mathrm{CO}_{2}\), and \(\mathrm{N}_{2}\). If the thrust of the rocket depends on the volume of the products produced, which of the substituted hydrazines produces a greater thrust per gram total mass of oxidizer plus fuel? (Assume that both fuels generate the same temperature and that \(\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) is formed.)

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 6 A element in each: (a) selenous acid, (b) potassium hydrogen sulfite, (c) hydrogen telluride, (d) carbon disulfide, (e) calcium sulfate.

Complete and balance the following equations: (a) \(\mathrm{CO}_{2}(g)+\mathrm{OH}^{-}(a q) \longrightarrow\) (b) \(\mathrm{NaHCO}_{3}(s)+\mathrm{H}^{+}(a q) \longrightarrow\) (c) \(\mathrm{CaO}(s)+\mathrm{C}(s) \stackrel{\Delta}{\longrightarrow}\) (d) \(\mathrm{C}(s)+\mathrm{H}_{2} \mathrm{O}(g) \stackrel{\Delta}{\longrightarrow}\) (e) \(\mathrm{CuO}(s)+\mathrm{CO}(g) \longrightarrow\)

Explain the following observations: (a) for a given oxidation state, the acid strength of the oxyacid in aqueous solution decreases in the order chlorine \(>\) bromine \(>\) iodine. (b) Hydrofluoric acid cannot be stored in glass bottles. (c) HI cannot be prepared by treating NaI with sulfuric acid. (d) The interhalogen \(\mathrm{ICl}_{3}\) is known, but \(\mathrm{BrCl}_{3}\) is not.

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (d) \(\mathrm{NH}_{3}(a q)+\mathrm{H}^{+}(a q) \longrightarrow\) (e) \(\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow\)
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