10: Electrochemistry (Supplementary)
- Page ID
- 475297
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)Electrochemistry is the study of chemical processes that cause electrons to move. This movement of electrons is called electricity, which can be generated by movements of electrons from one element to another in a reaction known as an oxidation-reduction ("redox") reaction. In this chapter, the married subjects of thermodynamics and equilibrium will be expanded to a troika by including electrochemistry. The ability to "steal" and "donate" electrons will be introduced and discussed in the context of free energies, enthalpies, entropies and equilibrium constants.
- 10.1: Electrochemical Cells
- This page explains the operation of galvanic cells, which convert spontaneous redox reactions into electrical energy, contrasting with electrolytic cells that require external energy. It details how a zinc strip oxidizes while a copper strip reduces, with a salt bridge maintaining charge neutrality. The text covers measuring voltage, the importance of balancing half-reactions, the function of electrodes, and provides examples of constructing cell diagrams to represent electrochemical cells.
- 10.2: The Gibbs Free Energy and Cell Voltage
- This page explores the interplay between cell potential, equilibrium constants, and spontaneity in electrochemical reactions. It details how concentration changes influence reaction spontaneity, energy conversion in electrochemical cells, and introduces key concepts from Michael Faraday. The text discusses calculating standard free energy change (\(ΔG^o\)) and demonstrates its implications for reaction spontaneity.
- 10.3: Concentration Effects and the Nernst Equation
- This page explains the relationship between cell potentials and Gibbs energy using the Nernst equation, particularly under non-standard conditions. It details the behavior of electrochemical cells, specifically how cell voltage decreases as reactions progress and reaches equilibrium. Examples, including the calculation of solubility products for compounds like AgCl and PbSO4 using electrochemical methods, illustrate the application of the Nernst equation.
- 10.4: Batteries and Fuel Cells
- This page provides an overview of various types of batteries and fuel cells, highlighting their chemical reactions, applications, and pros/cons. It details primary and secondary batteries like the Leclanché dry cell, alkaline, silver, mercury, lithium–iodine, NiCad, NiMH, and lead-acid batteries, discussing their efficiencies and drawbacks, including environmental concerns.
- 10.5: Corrosion and Its Prevention
- This page discusses corrosion, primarily affecting metals like iron through oxidation, leading to deterioration. It highlights rust formation due to iron's reaction with water and oxygen, and preventive methods such as coatings of noble metals and cathodic protection, which uses more anodic metals like zinc. Techniques to combat corrosion include protective coatings and sacrificial anodes, such as attaching zinc to galvanized steel or boats to protect against corrosion in various environments.
- 10.6: Electrolysis - A Deeper Look
- This page covers electrolysis, contrasting it with galvanic cells, and explains the functioning of electrolytic cells, particularly in producing aluminum and sodium. It discusses how electronegativity affects oxidation-reduction reactions during electrolysis, and the applications like electroplating.