23.8: Electrolytic Cells
- Page ID
- 53975
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)In 1989, two scientists announced that they had achieved "cold fusion", the process of fusing together elements at essentially room temperature to achieve energy production. The hypothesis was that the fusion would produce more energy than was required to cause the process to occur. Their process involved the electrolysis of heavy water (water molecules containing some deuterium instead of normal hydrogen) on a palladium electrode. The experiments could not be reproduced and their scientific reputations were pretty well shot. However, in more recent years, both industry and government researchers are taking another look at this process. The device illustrated above is part of a government project, and NASA is completing some studies on the topic as well. Cold fusion may not be so "cold" after all.
Electrolytic Cells
A voltaic cell uses a spontaneous redox reaction to generate an electric current. It is also possible to do the opposite. When an external source of direct current is applied to an electrochemical cell, a reaction that is normally nonspontaneous can be made to proceed. Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis is responsible for the appearance of many everyday objects, such as gold-plated or silver-plated jewelry, and chrome-plated car bumpers.
An electrolytic cell is the apparatus used for carrying out an electrolysis reaction. In an electrolytic cell, electric current is applied to provide a source of electrons for driving the reaction in a nonspontaneous direction. In a voltaic cell, the reaction goes in a direction that releases electrons spontaneously. In an electrolytic cell, the input of electrons from an external source forces the reaction to go in the opposite direction.
The spontaneous direction for the reaction between \(\ce{Zn}\) and \(\ce{Cu}\) is for the \(\ce{Zn}\) metal to be oxidized to \(\ce{Zn^{2+}}\) ions, while the \(\ce{Cu^{2+}}\) ions are reduced to \(\ce{Cu}\) metal. This makes the zinc electrode the anode and the copper electrode the cathode. When the same half-cells are connected to a battery via the external wire, the reaction is forced to run in the opposite direction. The zinc electrode is now the cathode and the copper electrode is the anode.
\[\begin{array}{lll} \text{oxidation (anode)} & \ce{Cu} \left( s \right) \rightarrow \ce{Cu^{2+}} \left( aq \right) + 2 \ce{e^-} & E^0 = -0.34 \: \text{V} \\ \text{reduction (cathode)} & \ce{Zn^{2+}} \left( aq \right) + 2 \ce{e^-} \rightarrow \ce{Zn} \left( s \right) & E^0 = -0.76 \: \text{V} \\ \hline \text{overall reaction} & \ce{Cu} \left( s \right) + \ce{Zn^{2+}} \left( aq \right) \rightarrow \ce{Cu^{2+}} \left( aq \right) + \ce{Zn} \left( s \right) & E^0_\text{cell} = -1.10 \: \text{V} \end{array}\nonumber \]
The standard cell potential is negative, indicating a nonspontaneous reaction. The battery must be capable of delivering at least \(1.10 \: \text{V}\) of direct current in order for the reaction to occur. Another difference between a voltaic cell and an electrolytic cell is the signs of the electrodes. In a voltaic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is negative and the cathode is positive. In an electrolytic cell, the anode is positive because it is connected to the positive terminal of the battery. The cathode is therefore negative. Electrons still flow through the cell from the anode to the cathode.
Summary
- The function of an electrolytic cell is described.
- Reactions illustrating electrolysis are given.