11: Electrochemistry
- Page ID
- 518125
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 11.0: Introduction
- The Economical Impact of Corrosion
- 11.1: Review of Redox Chemistry
- Redox reactions involve the transfer of electrons between species. In these reactions, one substance is oxidized (loses electrons) and another is reduced (gains electrons). Oxidation numbers are a helpful tool for identifying which atoms are oxidized or reduced and for determining whether a reaction is a redox process. To balance redox reactions, especially in aqueous solutions, the half-reaction method provides a systematic approach that ensures both atoms and charges are conserved.
- 11.2: Galvanic Cells
- Galvanic cells use a spontaneous redox reaction to generate electricity by separating oxidation and reduction into two half-cells connected by a salt bridge. Electrons flow through an external wire from the anode (oxidation) to the cathode (reduction), allowing the reaction to do work. Galvanic cells are described using cell notation and may use active or inert electrodes.
- 11.3: Cell Potentials
- The cell potential is a measure of the driving force for a redox reaction. Electron flow in a galvanic cell depends on the identity and concentrations of the reacting substances, and the temperature. The cell potential under standard conditions is called the standard cell potential. Standard reduction potentials can be used to calculate the standard cell potential and predict the direction of redox reactions under standard conditions.
- 11.4: Connecting Cell Potential, Free Energy, and Equilibrium
- Cell potential, free energy, and equilibrium are connected for redox reactions. The free energy change of a redox reaction represents the maximum electrical work a galvanic cell can perform on the surroundings. Under standard state conditions, a spontaneous redox reaction has a positive standard cell potential, a negative free energy change, and an equilibrium constant greater than 1.
- 11.5: Cell Potentials Beyond Standard Conditions
- The Nernst Equation allows us to calculate cell potentials under non-standard conditions and to predict the direction of redox reactions as concentrations change. It explains how concentration cells generate voltage and how galvanic cells can be used to measure ion concentrations in solution.
- 11.6: Corrosion
- Corrosion is the degradation of a metal through electrochemical processes. Iron corrodes in the presence of oxygen and water, and unlike many metals, the rust that forms does not create a protective layer. Corrosion can be prevented by eliminating factors that promote it or by using a more reactive metal that oxidizes in place of the iron.
- 11.7: Electrolysis
- Electrolysis uses electrical energy to drive nonspontaneous redox reactions, with oxidation at the anode and reduction at the cathode. Quantitative relationships link current, time, and mass of material deposited or consumed.
- 11.E: Exercises on Electrochemistry
- Working through the end-of-chapter problems is an important step in mastering electrochemistry. These questions help you apply core concepts such as oxidation–reduction, cell potentials, and the Nernst equation to a variety of systems.