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Chapter 17: Problem 11

Concept Review What is meant by a half-reaction?

Short Answer

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Answer: A half-reaction is a representation of a part of a redox reaction, where the oxidation and reduction processes are separated, and only the movement of electrons for one chemical species is considered. Half-reactions are significant in analyzing redox reactions as they help in identifying the species that are oxidized and reduced, determining the number of electrons transferred, and balancing redox reactions. They provide a clearer understanding of the electron transfer process, which is integral to redox chemistry.

Step by step solution

01

Definition of Half-reaction

A half-reaction is a representation of a part of a redox (oxidation-reduction) reaction, where we separate the overall reaction into two processes - the oxidation and the reduction. In a half-reaction, we only consider the movement of the electrons, either gaining or losing, for one chemical species. It's essential to note that half-reactions don't occur on their own, but they're helpful for analyzing and balancing redox reactions.
02

Examples of Half-reactions

To provide a clearer understanding, let's examine two half-reactions in the redox reaction between magnesium and oxygen: Mg -> Mg^2+ + 2e- (Oxidation half-reaction) 1/2 O_2 + 2e- -> O^2- (Reduction half-reaction) In the oxidation half-reaction, magnesium loses two electrons and forms Mg^2+. In the reduction half-reaction, each oxygen atom gains two electrons, converting it to the O^2- ion. Combining these half-reactions gives us the overall redox reaction: 2 Mg + O_2 -> 2 MgO
03

Significance of Half-reactions

Half-reactions play a crucial role in analyzing redox reactions, as they help us to: 1. Identify which species are oxidized and reduced. 2. Determine the number of electrons transferred. 3. Balance the redox reactions, including those involving acidic or basic solutions. By breaking redox reactions into half-reactions, we can better understand the electron transfer process, which is the central feature of redox chemistry.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Redox Reactions
Redox reactions are a type of chemical reaction that involves the transfer of electrons between two substances. This is an essential process for a vast array of chemical and biological systems. Redox reactions are characterized by two simultaneous sub-processes: oxidation and reduction. In these reactions, one substance loses electrons, while another gains them. This electron transfer can significantly change the chemical properties of the substances involved, leading to the formation of new compounds. A key feature of redox reactions is the involvement of half-reactions. These are simplified representations showing either the oxidation or reduction process separately. By examining half-reactions, it's possible to analyze and balance the overall redox equation more efficiently.
Oxidation
Oxidation is a part of redox reactions where a chemical species loses electrons. This process results in an increase in the oxidation state of the element. Losing electrons means that the substance undergoing oxidation becomes more positively charged. To visualize oxidation, imagine a piece of iron rusting in oxygen. The iron (Fe) atoms lose electrons to oxygen molecules, resulting in iron oxide or rust. In this process:
  • Iron is oxidized to form iron ions (Fe2+ or Fe3+).
  • The oxidation state of iron increases due to the loss of electrons.
In redox reactions, oxidation cannot happen independently; it always occurs alongside reduction. Thus, oxidation is paired with electron loss while another species is simultaneously gaining those electrons.
Reduction
Reduction is the complementary process to oxidation in a redox reaction. It involves the gain of electrons by a chemical species, leading to a decrease in the oxidation state. Gaining electrons generally makes the substance more negatively charged. A classic example of reduction can be observed when oxygen molecules (O2) gain electrons to form oxide ions (O2-). In this scenario:
  • Oxygen molecules are reduced as they gain electrons.
  • The overall charge on the oxygen atoms becomes more negative.
Like oxidation, reduction happens in the context of a redox reaction. As one element loses electrons, another must gain them. This duality ensures that the charge balance in the system is maintained, reflecting the conservation of charge principle.
Electron Transfer
Electron transfer is the fundamental process at the heart of redox reactions. As mentioned, redox reactions involve the shuffling of electrons between substances, driving both oxidation and reduction. The significance of electron transfer lies in its ability to change the chemical nature of the involved substances. During this process:
  • Electrons move from a donor (oxidized species) to an acceptor (reduced species).
  • This movement can create new chemical bonds or break existing ones, resulting in the formation of new compounds.
One illustrative application of electron transfer is in electrochemical reactions, such as those occurring in batteries. Here, controlled electron flow from the anode to the cathode generates electrical energy. Understanding how electrons transfer in redox reactions is crucial for designing chemical processes and technologies that rely on electron flow.

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Most popular questions from this chapter

Relating Battery Capacity to Quantities of Reactants One 12 -volt lead-acid battery has a higher ampere \cdot hour rating than another. Which of the following parameters are likely to be different for the two batteries? a. Individual cell potentials b. Anode half-reactions c. Total masses of electrode materials d. Number of cells e. Electrolyte composition f. Combined surface areas of their electrodes

The standard potential of the \(\mathrm{Zn} / \mathrm{Cu}^{2+}\) cell reaction $$ \mathrm{Zn}(s)+\mathrm{Cu}^{2+}(a q) \rightarrow \mathrm{Zn}^{2+}(a q)+\mathrm{Cu}(s) $$ is \(1.10 \mathrm{V}\). Would the potential of the cell differ from \(1.10 \mathrm{V}\) if the concentrations of both \(\mathrm{Cu}^{2+}\) and \(\mathrm{Zn}^{2+}\) were \(0.25 M ?\)

From the table of standard reduction potentials in Appendix 6 a. Select an oxidizing agent that will oxidize \(\mathrm{Cr}(s)\) to \(\mathrm{Cr}^{3+}(a q)\) but not \(\mathrm{Cd}(s)\) to \(\mathrm{Cd}^{2+}(a q)\) b. Select a reducing agent that will reduce Br \(_{2}(\ell)\) to \(\mathrm{Br}^{-}(a q)\) but not \(\mathrm{I}_{2}(s)\) to \(1^{-}(a q).\)

What is the value of \(E^{\circ}\) for the following reaction? $$ 2 \mathrm{AgCl}(s)+\mathrm{H}_{2}(g) \rightarrow 2 \mathrm{Ag}(s)+2 \mathrm{HCl}(a q) $$

The salts obtained from the evaporation of seawater can act as a source of halogens, principally \(\mathrm{Cl}_{2}\) and \(\mathrm{Br}_{2}\), through the electrolysis of the molten alkali metal halides. As the potential of the anode in an electrolytic cell is increased, which of these two halogens forms first?
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