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Chapter 17: Problem 39

What is a half-reaction? What does each of the two half-reactions that make up an overall process represent?

Short Answer

Expert verified
A half-reaction is a chemical reaction that shows either the oxidation or reduction process separately, involving the transfer of electrons between chemical species. In a redox (reduction-oxidation) reaction, the overall process is accomplished by combining two half-reactions: an oxidation half-reaction, which represents a chemical species losing electrons and increasing its oxidation state, and a reduction half-reaction, which represents a chemical species gaining electrons and decreasing its oxidation state. When combined, the electrons lost in the oxidation half-reaction equal the electrons gained in the reduction half-reaction, ensuring that the overall process is balanced with respect to both mass and charge.

Step by step solution

01

Definition of Half-Reaction

A half-reaction refers to a chemical reaction that shows either oxidation or reduction process separately. It involves the transfer of electrons between chemical species. Half-reactions are often used in redox (reduction-oxidation) reactions, which involve both oxidation and reduction taking place. In a redox reaction, one reactant is oxidized (loses electrons), and another reactant is reduced (gains electrons). So, overall redox reaction can be separated into two half-reactions: one for oxidation and the other for reduction.
02

Oxidation Half-Reaction

The oxidation half-reaction represents the process where a chemical species loses one or more electrons and increases its oxidation state. This species is known as the reducing agent, as it donates electrons to another species. A common example of an oxidation half-reaction is the oxidation of solid zinc to zinc ions in an aqueous solution: \[ Zn(s) \rightarrow Zn^{2+}(aq) + 2e^{-} \]
03

Reduction Half-Reaction

The reduction half-reaction represents the process where a chemical species gains one or more electrons and decreases its oxidation state. This species is known as the oxidizing agent, as it accepts electrons from another species. A common example of a reduction half-reaction is the reduction of hydrogen ions to hydrogen gas in an aqueous solution: \[ 2H^{+}(aq) + 2e^{-} \rightarrow H_{2}(g) \]
04

Connecting the Two Half-Reactions

The overall process in a redox reaction is accomplished by combining the oxidation and reduction half-reactions. When the half-reactions are added together, the electrons lost in the oxidation half-reaction equal the electrons gained in the reduction half-reaction. This ensures that the overall process is balanced with respect to both mass and charge. In the example given above, the two half-reactions can be combined to form an overall reaction as follows: \[ Zn(s) + 2H^{+}(aq) \rightarrow Zn^{2+}(aq) + H_{2}(g) \] In this overall process, solid zinc (reducing agent) loses 2 electrons through oxidation and hydrogen ions (oxidizing agent) gain 2 electrons through reduction.

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Most popular questions from this chapter

For each of the following oxidation-reduction reactions, identify which element is being oxidized and which is being reduced. a. \(2 \mathrm{Fe}(s)+3 \mathrm{F}_{2}(g) \rightarrow 2 \mathrm{FeF}_{3}(s)\) b. \(\mathrm{O}_{2}(g)+2 \mathrm{Cu}(s) \rightarrow 2 \mathrm{CuO}(s)\) c. \(\mathrm{F}_{2}(g)+2 \mathrm{KI}(a q) \rightarrow 2 \mathrm{KF}(a q)+\mathrm{I}_{2}(s)\) d. \(2 \mathrm{Al}(s)+3 \mathrm{H}_{2}(g) \rightarrow 2 \mathrm{AlH}_{3}(s)\)

Carbon compounds containing double bonds (such compounds are called alkenes) react readily with many other reagents. In each of the following reactions, identify which atoms are oxidized and which are reduced, and specify the oxidizing and reducing agents. a. \(\mathrm{CH}_{2}=\mathrm{CH}_{2}(g)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{ClCH}_{2}-\mathrm{CH}_{2} \mathrm{Cl}(l)\) b. \(\mathrm{CH}_{2}=\mathrm{CH}_{2}(g)+\mathrm{Br}_{2}(g) \rightarrow \mathrm{BrCH}_{2}-\mathrm{CH}_{2} \operatorname{Br}(l)\) c. \(\mathrm{CH}_{2}=\mathrm{CH}_{2}(g)+\mathrm{HBr}(g) \rightarrow \mathrm{CH}_{3}-\mathrm{CH}_{2} \mathrm{Br}(l)\) d. \(\mathrm{CH}_{2}=\mathrm{CH}_{2}(g)+\mathrm{H}_{2}(g) \rightarrow \mathrm{CH}_{3}-\mathrm{CH}_{3}(g)\)

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. \(\mathrm{ClO}_{4}^{-}(a q)+\mathrm{Cl}^{-}(a q) \rightarrow \mathrm{ClO}_{3}^{-}(a q)+\mathrm{Cl}_{2}(g)\)

Balance each of the following half-reactions, which take place in acidic solution. a. \(\mathrm{HBr}(a q) \rightarrow \mathrm{Br}^{-}(a q)+\mathrm{H}_{2}(g)\) b. \(\mathrm{ClO}_{4}^{-}(a q) \rightarrow \mathrm{ClO}_{3}^{-}(a q)\) c. \(\operatorname{PbSO}_{4}(s) \rightarrow \operatorname{Pb}(s)+\operatorname{HSO}_{4}^{-}(a q)\) d. \(\mathrm{HNO}_{2}(a q) \rightarrow \mathrm{NO}(g)\)

Consider the oxidation-reduction reaction \(\mathrm{Mg}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Mg}^{2+}(a q)+\mathrm{Cu}(s)\) Sketch a galvanic cell that uses this reaction. Which metal ion is reduced? Which metal is oxidized? What half-reaction takes place at the anode in the cell? What half-reaction takes place at the cathode?
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