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Chapter 17: Problem 27

Does an oxidizing agent donate or accept electrons? Does a reducing agent donate or accept electrons?

Short Answer

Expert verified
In a redox reaction, an oxidizing agent accepts electrons and causes another species to get oxidized, while a reducing agent donates electrons and causes another species to get reduced.

Step by step solution

01

Introduction of redox reactions

A redox reaction, short for reduction-oxidation reaction, involves the transfer of electrons between two species. In a redox reaction, one species is reduced (gains electrons) and the other is oxidized (loses electrons).
02

Oxidizing agent

An oxidizing agent is a substance that readily accepts electrons from other species in a redox reaction, causing the other species to be oxidized.
03

Oxidizing agent's role in a redox reaction

An oxidizing agent does the following in a redox reaction: 1. Accepts electrons 2. Causes another species to get oxidized (lose electrons)
04

Reducing agent

A reducing agent is a substance that easily donates electrons to other species in a redox reaction, causing the other species to be reduced.
05

Reducing agent's role in a redox reaction

A reducing agent does the following in a redox reaction: 1. Donates electrons 2. Causes another species to get reduced (gain electrons) To summarize, an oxidizing agent accepts electrons and a reducing agent donates electrons in a redox reaction.

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Most popular questions from this chapter

For each of the following oxidation-reduction reactions of metals with nonmetals, identify which element is oxidized and which is reduced. a. \(4 \mathrm{Na}(s)+\mathrm{O}_{2}(g) \rightarrow 2 \mathrm{Na}_{2} \mathrm{O}(s)\) b. \(\operatorname{Fe}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \rightarrow \mathrm{FeSO}_{4}(a q)+\mathrm{H}_{2}(g)\) c. \(2 \mathrm{Al}_{2} \mathrm{O}_{3}(s) \rightarrow 4 \mathrm{Al}(s)+3 \mathrm{O}_{2}(g)\) d. \(3 \mathrm{Mg}(s)+\mathrm{N}_{2}(g) \rightarrow \mathrm{Mg}_{3} \mathrm{N}_{2}(s)\)

In order to obtain useful electrical energy from an oxidation-reduction process, we must set up the reaction in such a way that the oxidation half- reaction and the reduction half-reaction are physically __________ one another.

Balance each of the following oxidation-reduction reactions, which take place in acidic solution, by using the "half-reaction" method. a. \(\mathrm{Mg}(s)+\mathrm{Hg}^{2+}(a q) \rightarrow \mathrm{Mg}^{2+}(a q)+\mathrm{Hg}_{2}^{2+}(a q)\) b. \(\mathrm{NO}_{3}^{-}(a q)+\mathrm{Br}^{-}(a q) \rightarrow \mathrm{NO}(g)+\mathrm{Br}_{2}(l)\) c. \(\mathrm{Ni}(s)+\mathrm{NO}_{3}^{-}(a q) \rightarrow \mathrm{Ni}^{2+}(a q)+\mathrm{NO}_{2}(g)\) d. \(\mathrm{ClO}_{4}^{-}(a q)+\mathrm{Cl}^{-}(a q) \rightarrow \mathrm{ClO}_{3}^{-}(a q)+\mathrm{Cl}_{2}(g)\)

Why must the number of electrons lost in the oxidation equal the number of electrons gained in the reduction? Is it possible to have "leftover" electrons in a reaction?

Balance each of the following half-reactions. a. \(\mathrm{N}_{2}(g) \rightarrow \mathrm{N}^{3-}(s)\) b. \(\mathrm{O}_{2}^{2-}(a q) \rightarrow \mathrm{O}_{2}(g)\) c. \(\mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(a q)\) d. \(\mathrm{F}_{2}(g) \rightarrow \mathrm{F}^{-}(a q)\)
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