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Chapter 21: Problem 2

Why must an electrochemical process involve a redox reaction?

Short Answer

Expert verified
Electrochemical processes require redox reactions because they involve the transfer of electrons.

Step by step solution

01

Understand Electrochemical Processes

An electrochemical process is a chemical reaction that involves the transfer of electrons between two species. This transfer of electrons generates electrical energy or consumes electrical energy to drive the reaction.
02

Define Redox Reactions

Redox reactions, or reduction-oxidation reactions, are chemical reactions involving the transfer of electrons between two substances. One substance loses electrons (oxidation), while the other gains electrons (reduction).
03

Relationship Between Electrochemistry and Redox Reactions

For an electrochemical process to occur, there must be a transfer of electrons, which is the essence of redox reactions. Without redox reactions, the flow of electrons necessary for generating or consuming electrical energy cannot happen. Thus, an electrochemical process must involve a redox reaction.
04

Conclusion

In summary, electrochemical processes require the movement of electrons to function, and this movement is facilitated by redox reactions, where oxidation and reduction occur simultaneously.

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Key Concepts

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

Redox Reactions
Redox reactions, or reduction-oxidation reactions, are fundamental chemical processes where electrons are transferred between substances. In a redox reaction, one substance loses electrons, a process known as oxidation. Simultaneously, the other substance gains those electrons, a process called reduction. This exchange is crucial as it drives many chemical and biological processes.
Without redox reactions, essential mechanisms such as respiration and photosynthesis in living organisms could not occur. The 'redox' term itself is a combination of reduction and oxidation, reflecting the inseparable nature of these two processes.
Electron Transfer
Electron transfer is at the heart of electrochemical processes. It involves the movement of electrons from one atom or molecule to another. This transfer is essential because it enables the flow of electrical energy.
The movement of electrons can be spontaneous, releasing energy, or it can require external energy to proceed. In batteries, for example, electrons move from one electrode to another, generating electric current that powers devices.
Understanding electron transfer helps in grasping how various redox reactions power our everyday lives and industrial applications. It's not just about atoms and molecules; it's about channelling their interactions to produce energy.
Oxidation and Reduction
In any redox reaction, oxidation and reduction always occur together. They are complementary processes that involve either the loss or gain of electrons.
  • Oxidation: This is when a substance loses electrons. It increases the oxidation state of the substance. For example, in the reaction between zinc and copper sulfate, zinc gets oxidized as it loses electrons and forms zinc ions.

  • Reduction: This is when a substance gains electrons. It decreases the oxidation state of the substance. Continuing the zinc and copper sulfate example, copper ions gain the electrons lost by zinc, thereby reducing to form metallic copper.
When you see one happening, the other is occurring too. Both processes are indispensable in driving the redox reactions necessary for electrochemical processes.

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Most popular questions from this chapter

How many grams of aluminum can form by passing 305 C through an electrolytic cell containing a molten aluminum salt?

What does a negative \(E_{\mathrm{cell}}^{\circ}\) indicate about a redox reaction? What does it indicate about the reverse reaction?

To examine the effect of ion removal on cell voltage, a chemist constructs two voltaic cells, each with a standard hydrogen electrode in one compartment. One cell also contains a \(\mathrm{Pb} / \mathrm{Pb}^{2+}\) half-cell; the other contains a \(\mathrm{Cu} / \mathrm{Cu}^{2+}\) half-cell. (a) What is \(E^{\circ}\) of each cell at \(298 \mathrm{~K} ?\) (b) Which electrode in each cell is negative? (c) When \(\mathrm{Na}_{2} \mathrm{~S}\) solution is added to the \(\mathrm{Pb}^{2+}\) electrolyte, solid \(\mathrm{PbS}\) forms. What happens to the cell voltage? (d) When sufficient \(\mathrm{Na}_{2} \mathrm{~S}\) is added to the \(\mathrm{Cu}^{2+}\) electrolyte, CuS forms and \(\left[\mathrm{Cu}^{2+}\right]\) drops to \(1 \times 10^{-16} \mathrm{M} .\) Find the cell voltage.

Like any piece of apparatus, an electrolytic cell operates at less than \(100 \%\) efficiency. A cell depositing \(\mathrm{Cu}\) from a \(\mathrm{Cu}^{2+}\) bath operates for \(10 \mathrm{~h}\) with an average current of 5.8 A. If \(53.4 \mathrm{~g}\) of copper is deposited, at what efficiency is the cell operating?

Commercial electrolytic cells for producing aluminum operate at \(5.0 \mathrm{~V}\) and \(100,000 \mathrm{~A}\). (a) How long does it take to produce exactly 1 metric ton ( \(1000 \mathrm{~kg}\) ) of aluminum? (b) How much electrical power (in kilowatt-hours, \(\mathrm{kW} \cdot \mathrm{h}\) ) is used \(\left(1 \mathrm{~W}=1 \mathrm{~J} / \mathrm{s} ; 1 \mathrm{~kW} \cdot \mathrm{h}=3.6 \times 10^{3} \mathrm{~kJ}\right) ?\) (c) If electricity costs \(\$ 0.123\) per \(\mathrm{kW} \cdot \mathrm{h}\) and cell efficiency is \(90 . \%\) what is the cost of electricity to produce exactly 1 lb of aluminum?
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