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Chapter 20: Problem 124

Construct a concept map illustrating the relationship between batteries and electrochemical ideas.

Short Answer

Expert verified
A concept map can illustrate that electrochemistry is a field that studies the movement of electrons and includes topics like redox reactions and electrochemical cells. Batteries, which utilize these principles, convert chemical energy to electrical energy. The whole process involves the flow of electrons from the anode to the cathode, aided by the movement of cations and anions.

Step by step solution

01

Identifying Main Concepts

The first step towards creating a concept map is identifying the main concepts that must be included: Batteries, Electrochemistry, Redox reactions, Cations and Anions, Electrochemical cells, Anode, Cathode, and Electron flow.
02

Establish Relationship between Concepts

Next, describe how these concepts are connected. For example, Electrochemistry is the overarching concept studying chemical processes that cause electrons to move. This involves Redox reactions which are reactions where oxidation and reduction occur simultaneously. Batteries are a practical application of these principles. They convert chemical energy to electrical energy through a process that occurs in an electrochemical cell.
03

Arrange the Concepts

Now, arrange the concepts in the concept map. Start with the broad topic like 'Electrochemistry' and then branch out to smaller, detailed concepts such as 'Redox reactions', 'Electrochemical cells' and 'Batteries' and explain their relations.
04

Draw the Connections

Finally, draw the connections between the concepts. These can be shown as arrows or lines between concepts. Each connection should be labelled with linking phrases or words that describe the relationship between the concepts. For example, an arrow pointing from 'Redox reactions' to 'Batteries' could be labelled with 'occur in'.

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Key Concepts

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

Batteries
Batteries are fascinating devices that power many of the tools we use every day, from remote controls to electric cars. Essentially, a battery is a self-contained system that converts chemical energy into electrical energy through reactions within its cells. Inside each battery, there are several electrochemical cells working together. Each cell consists of two electrodes (an anode and a cathode) and an electrolyte that facilitates the movement of ions.
  • The chemical reactions occurring at these electrodes produce a flow of electrons, creating electricity.
  • Batteries come in various types, including primary (non-rechargeable) and secondary (rechargeable) batteries.
  • The electricity produced by batteries powers circuits and provides energy for electronic devices.
Understanding battery structure and function is crucial for unlocking the power of electrochemistry.
Redox reactions
Redox reactions, short for oxidation-reduction reactions, are at the heart of electrochemistry. These reactions involve the transfer of electrons between substances. In every redox reaction, one species is oxidized, meaning it loses electrons, and another is reduced, meaning it gains electrons. This electron movement is essential for many chemical processes.
Redox reactions play a critical role in batteries. When a battery operates, redox reactions occur at the electrodes: one electrode donates electrons (oxidation), and the other accepts them (reduction). This flow of electrons between the electrodes generates an electric current.
  • Oxidation happens at the anode.
  • Reduction occurs at the cathode.
Understanding redox reactions is key to comprehending how batteries transform energy effectively.
Electrochemical cells
Electrochemical cells are the foundational units of a battery. Each cell is a small, intricate component that converts chemical energy into electrical energy through redox reactions. There are two main types of electrochemical cells:
  • Galvanic (or Voltaic) Cells: These generate electricity spontaneously through redox reactions. They are commonly used in batteries.
  • Electrolytic Cells: Electricity is used to drive non-spontaneous reactions in these cells, such as electroplating.
In a typical electrochemical cell, the anode is where oxidation occurs, while the cathode is where reduction takes place. These electrodes are immersed in an electrolyte solution that allows ions to move, facilitating the chemical reactions needed for electricity production.
Learning about electrochemical cells helps understand how batteries store energy and release it to do work.
Cations and Anions
At the molecular level, cations and anions are charged ions that play a crucial role in ensuring that electrochemical processes proceed smoothly.
  • Cations: Positively charged ions that move towards the cathode during electrochemical reactions.
  • Anions: Negatively charged ions that move towards the anode.
In the context of a battery, the movement of cations and anions is vital for maintaining electrical neutrality as electrons flow through the external circuit. The electrolyte in the battery is full of these ions, enabling the internal charge transfer necessary for sustained electron flow.
Understanding cations and anions is critical to grasping how batteries function and the subtle interplay of chemicals within electrochemical cells.

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Most popular questions from this chapter

Show that for a combination of half-cell reactions that produce a standard reduction potential for a half-cell that is not directly observable, the standard reduction potential is $$E^{\circ}=\frac{\sum n_{i} E_{i}^{\circ}}{\sum n_{i}}$$ where \(n_{i}\) is the number of electrons in each half-reaction of potential \(E_{i}^{\circ} .\) Use the following half-reactions: $$ \begin{array}{c} \mathrm{H}_{5} \mathrm{IO}_{6}(\mathrm{aq})+\mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{IO}_{3}^{-}(\mathrm{aq})+ \\ 3 \mathrm{H}_{2} \mathrm{O}(1) \quad E^{\circ}=1.60 \mathrm{V} \\ \mathrm{IO}_{3}^{-}(\mathrm{aq})+6 \mathrm{H}^{+}(\mathrm{aq})+5 \mathrm{e}^{-} \longrightarrow \frac{1}{2} \mathrm{I}_{2}(\mathrm{s})+3 \mathrm{H}_{2} \mathrm{O}(1) \\ E^{\circ}=1.19 \mathrm{V} \\ 2 \mathrm{HIO}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{I}_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(1) \\ E^{\circ}=1.45 \mathrm{V} \\ \mathrm{I}_{2}(\mathrm{s})+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{I}^{-}(\mathrm{aq}) \quad \quad E^{\circ}=0.535 \mathrm{V} \end{array} $$ Calculate the standard reduction potential for $$ \mathrm{H}_{6} \mathrm{IO}_{6}+5 \mathrm{H}^{+}+2 \mathrm{I}^{-}+3 \mathrm{e}^{-} \longrightarrow $$ $$ \frac{1}{2} \mathrm{I}_{2}+4 \mathrm{H}_{2} \mathrm{O}=2 \mathrm{HIO} $$

When deciding whether a particular reaction corresponds to a cell with a positive standard cell potential, which of the following thermodynamic properties would you use to get your answer without performing any calculations? Which would you not use? Explain. (a) \(\Delta G^{\circ} ;\) (b) \(\Delta S^{\circ} ;\) (c) \(\Delta H^{\circ} ;\) (d) \(\Delta U^{\circ} ;\) (e) \(K\).

An aqueous solution of \(\mathrm{K}_{2} \mathrm{SO}_{4}\) is electrolyzed by means of Pt electrodes. (a) Which of the following gases should form at the anode: \(\mathrm{O}_{2}, \mathrm{H}_{2}, \mathrm{SO}_{2}, \mathrm{SO}_{3} ?\) Explain. (b) What product should form at the cathode? Explain. (c) What is the minimum voltage required? Why is the actual voltage needed likely to be higher than this value?

For the reaction \(2 \mathrm{Cu}^{+}(\mathrm{aq})+\mathrm{Sn}^{4+}(\mathrm{aq}) \longrightarrow\) \(2 \mathrm{Cu}^{2+}(\mathrm{aq})+\mathrm{Sn}^{2+}(\mathrm{aq}), E_{\mathrm{cell}}^{\circ}=-0.0050 \mathrm{V}\) (a) can a solution be prepared at \(298 \mathrm{K}\) that is \(0.500 \mathrm{M}\) in each of the four ions? (b) If not, in which direction will a reaction occur?

Given these half-reactions and associated standard reduction potentials, answer the questions that follow: $$\begin{aligned} &\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Zn}(\mathrm{s})+4 \mathrm{NH}_{3}(\mathrm{aq})\\\ &E^{\circ}=-1.015 \mathrm{V} \end{aligned}$$ $$\begin{array}{c} \mathrm{Ti}^{3+}(\mathrm{aq})+\mathrm{e}^{-} \longrightarrow \mathrm{Ti}^{2+}(\mathrm{aq}) \\ E^{\circ}=-0.37 \mathrm{V} \end{array}$$ $$\begin{aligned} &\mathrm{VO}^{2+}(\mathrm{aq})+2 \mathrm{H}^{+}(\mathrm{aq})+\mathrm{e}^{-} \longrightarrow \mathrm{V}^{3+}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{I})\\\ &E^{\circ}=0.340 \mathrm{V} \end{aligned}$$ $$\begin{array}{r} \mathrm{Sn}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Sn}(\mathrm{aq}) \\ E^{\circ}=-0.14 \mathrm{V} \end{array}$$ (a) Determine which pair of half-cell reactions leads to a cell reaction with the largest positive cell potential, and calculate its value. Which couple is at the anode and which is at the cathode? (b) Determine which pair of these half-cell reactions leads to the cell with the smallest positive cell potential, and calculate its value. Which couple is at the anode and which is at the cathode?
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