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<title>Standard electrode potential</title>

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<span id="openzim-page-title" class="mw-page-title-main"><span class="mw-page-title-main">Standard electrode potential</span></span>

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<p>In <a href="Electrochemistry" title="Electrochemistry">electrochemistry</a>, <b>standard electrode potential</b> <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle E^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle E^{\ominus }}</annotation>

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</math></span><img src="./d79e73b6c90d0170701b9e47deb3bda880d90584.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.338ex; width:3.305ex; height:2.509ex;" alt="{\displaystyle E^{\ominus }}" loading="lazy"></span>, or <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle E_{red}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle E_{red}^{\ominus }}</annotation>

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</math></span><img src="./2eca8e3be4f894ae5145e50114c550b2b2df77fa.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:4.315ex; height:3.176ex;" alt="{\displaystyle E_{red}^{\ominus }}" loading="lazy"></span>, is the <a href="Electrode_potential" title="Electrode potential">electrode potential</a> (a measure of the reducing power of any element or compound) which the IUPAC "Gold Book" defines as <i>"the value of the standard <a href="Electromotive_force" title="Electromotive force">emf</a> (<a href="Electromotive_force" title="Electromotive force">electromotive force</a>) of a cell in which molecular hydrogen under <a href="Standard_pressure" class="mw-redirect" title="Standard pressure">standard pressure</a> is oxidized to solvated protons at the left-hand electrode"</i>.<sup id="cite_ref-1" class="reference"><a href="#cite_note-1"><span class="cite-bracket">[</span>1<span class="cite-bracket">]</span></a></sup>

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<div class="mw-heading mw-heading2"><h2 id="Background">Background</h2></div>

<p>The basis for an <a href="Electrochemical_cell" title="Electrochemical cell">electrochemical cell</a>, such as the <a href="Galvanic_cell" title="Galvanic cell">galvanic cell</a>, is always a <a href="Redox_reaction" class="mw-redirect" title="Redox reaction">redox reaction</a> which can be broken down into two <a href="Half-reaction" title="Half-reaction">half-reactions</a>: <a href="Oxidation" class="mw-redirect" title="Oxidation">oxidation</a> at anode (loss of electron) and <a href="Redox" title="Redox">reduction</a> at cathode (gain of electron). <a href="Electricity" title="Electricity">Electricity</a> is produced due to the difference of <a href="Electric_potential" title="Electric potential">electric potential</a> between the individual potentials of the two metal <a href="Electrodes" class="mw-redirect" title="Electrodes">electrodes</a> with respect to the <a href="Electrolyte" title="Electrolyte">electrolyte</a>.

</p><p>Although the overall potential of a cell can be measured, there is no simple way to accurately measure the <a href="Absolute_electrode_potential" title="Absolute electrode potential">electrode/electrolyte potentials</a> in isolation. The electric potential also varies with temperature, concentration and pressure. Since the oxidation potential of a half-reaction is the negative of the reduction potential in a redox reaction, it is sufficient to calculate either one of the potentials. Therefore, standard electrode potential is commonly written as standard reduction potential.

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<div class="mw-heading mw-heading2"><h2 id="Calculation">Calculation</h2></div>

<p>The <a href="Galvanic_cell" title="Galvanic cell">galvanic cell</a> potential results from the voltage difference of a <i>pair</i> of electrodes. It is not possible to measure an absolute value for each electrode separately. However, the potential of a reference electrode, <a href="Standard_hydrogen_electrode" title="Standard hydrogen electrode">standard hydrogen electrode</a> (SHE), is defined as to 0.00 V. An electrode with unknown electrode potential can be paired with either the standard hydrogen electrode, or another electrode whose potential has already been measured, to determine its "absolute" potential.

</p><p>Since the electrode potentials are conventionally defined as reduction potentials, the sign of the potential for the metal electrode being oxidized must be reversed when calculating the overall cell potential. The electrode potentials are independent of the number of electrons transferred —they are expressed in volts, which measure energy per electron transferred—and so the two electrode potentials can be simply combined to give the overall <i>cell</i> potential even if different numbers of electrons are involved in the two electrode reactions.

</p><p>For practical measurements, the electrode in question is connected to the positive terminal of the <a href="Electrometer" title="Electrometer">electrometer</a>, while the standard hydrogen electrode is connected to the negative terminal.<sup id="cite_ref-2" class="reference"><a href="#cite_note-2"><span class="cite-bracket">[</span>2<span class="cite-bracket">]</span></a></sup>

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<div class="mw-heading mw-heading2"><h2 id="Reversible_electrode">Reversible electrode</h2></div>

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</style><div role="note" class="hatnote navigation-not-searchable">See also: <a href="Reversible_hydrogen_electrode" title="Reversible hydrogen electrode">Reversible hydrogen electrode</a></div>

<p>A reversible electrode is an electrode that owes its potential to <a href="Reversible_reaction" title="Reversible reaction">changes of a reversible nature</a>. A first condition to be fulfilled is that the system is close to the <a href="Chemical_equilibrium" title="Chemical equilibrium">chemical equilibrium</a>. A second set of conditions is that the system is submitted to very small solicitations spread on a sufficient period of time so, that the chemical equilibrium conditions nearly always prevail. In theory, it is very difficult to experimentally achieve reversible conditions because any perturbation imposed to a system near equilibrium in a finite time forces it out of equilibrium. However, if the solicitations exerted on the system are sufficiently small and applied slowly, one can consider an electrode to be reversible. By nature, electrode reversibility depends on the experimental conditions and the way the electrode is operated. For example, electrodes used in electroplating are operated with a high over-potential to force the reduction of a given metal cation to be deposited onto a metallic surface to be protected. Such a system is far from equilibrium and continuously submitted to important and constant changes in a short period of time

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<div class="mw-heading mw-heading2"><h2 id="Standard_reduction_potential_table">Standard reduction potential table</h2></div>

<div role="note" class="hatnote navigation-not-searchable">Main article: <a href="Standard_electrode_potential_(data_page)" title="Standard electrode potential (data page)">Standard electrode potential (data page)</a></div>

<div role="note" class="hatnote navigation-not-searchable">See also: <a href="Table_of_standard_reduction_potentials_for_half-reactions_important_in_biochemistry" title="Table of standard reduction potentials for half-reactions important in biochemistry">Table of standard reduction potentials for half-reactions important in biochemistry</a></div>

<p>The larger the value of the standard reduction potential, the easier it is for the element to be reduced (gain <a href="Electron" title="Electron">electrons</a>); in other words, they are better <a href="Oxidizing_agent" title="Oxidizing agent">oxidizing agents</a>.

</p><p>For example, F<sub>2</sub> has a standard reduction potential of +2.87 V and Li<sup>+</sup> has −3.05 V:

</p>

<p>The highly positive standard reduction potential of F<sub>2</sub> means it is reduced easily and is therefore a good oxidizing agent. In contrast, the greatly negative standard reduction potential of Li<sup>+</sup> indicates that it is not easily reduced. Instead, Li<sub>(<i>s</i>)</sub> would rather undergo oxidation (hence it is a good <a href="Reducing_agent" title="Reducing agent">reducing agent</a>).

</p><p>Zn<sup>2+</sup> has a standard reduction potential of −0.76 V and thus can be oxidized by any other electrode whose standard reduction potential is greater than −0.76 V (e.g., H<sup>+</sup> (0 V), Cu<sup>2+</sup> (0.34 V), F<sub>2</sub> (2.87 V)) and can be <a href="Redox" title="Redox">reduced</a> by any electrode with standard reduction potential less than −0.76 V (e.g. H<sub>2</sub> (−2.23 V), Na<sup>+</sup> (−2.71 V), Li<sup>+</sup> (−3.05 V)).

</p><p>In a galvanic cell, where a <a href="Spontaneous_process" title="Spontaneous process">spontaneous</a> redox reaction drives the cell to produce an electric potential, <a href="Gibbs_free_energy" title="Gibbs free energy">Gibbs free energy</a> <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle \Delta G^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle \Delta G^{\ominus }}</annotation>

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</math></span><img src="./783df781bd2cde06d7956470e952c27226a4e203.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -0.338ex; width:5.273ex; height:2.509ex;" alt="{\displaystyle \Delta G^{\ominus }}" loading="lazy"></span> must be negative, in accordance with the following equation:

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<annotation encoding="application/x-tex">{\displaystyle \Delta G_{cell}^{\ominus }=-nFE_{cell}^{\ominus }}</annotation>

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</math></span><img src="./9a8156d96f12d2995c201907cff642f3d925e89b.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:18.902ex; height:3.176ex;" alt="{\displaystyle \Delta G_{cell}^{\ominus }=-nFE_{cell}^{\ominus }}" loading="lazy"></span>      (unit: Joule = Coulomb × Volt)</dd></dl>

<p>where <span class="texhtml mvar" style="font-style:italic;">n</span> is number of <a href="Mole_(unit)" title="Mole (unit)">moles</a> of electrons per mole of products and <span class="texhtml mvar" style="font-style:italic;">F</span> is the <a href="Faraday_constant" title="Faraday constant">Faraday constant</a>, <span class="nowrap">~ 96 485 C/mol</span>.

</p><p>As such, the following rules apply:

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<annotation encoding="application/x-tex">{\displaystyle E_{cell}^{\ominus }}</annotation>

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</math></span><img src="./53586c734844c240fa1837ba007fedf940cf37a3.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:4.406ex; height:3.176ex;" alt="{\displaystyle E_{cell}^{\ominus }}" loading="lazy"></span> > 0, then the process is spontaneous (<a href="Galvanic_cell" title="Galvanic cell">galvanic cell</a>): <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle \Delta G_{cell}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle \Delta G_{cell}^{\ominus }}</annotation>

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</math></span><img src="./2b000458eb6269178d0ef5e50f33d5cc4c89d385.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:6.453ex; height:3.176ex;" alt="{\displaystyle \Delta G_{cell}^{\ominus }}" loading="lazy"></span> < 0, and energy is liberated.</dd></dl>

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<annotation encoding="application/x-tex">{\displaystyle E_{cell}^{\ominus }}</annotation>

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</math></span><img src="./53586c734844c240fa1837ba007fedf940cf37a3.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:4.406ex; height:3.176ex;" alt="{\displaystyle E_{cell}^{\ominus }}" loading="lazy"></span> < 0, then the process is non-spontaneous (<a href="Electrolytic_cell" title="Electrolytic cell">electrolytic cell</a>): <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle \Delta G_{cell}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle \Delta G_{cell}^{\ominus }}</annotation>

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</math></span><img src="./2b000458eb6269178d0ef5e50f33d5cc4c89d385.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:6.453ex; height:3.176ex;" alt="{\displaystyle \Delta G_{cell}^{\ominus }}" loading="lazy"></span> > 0, and energy is consumed.</dd></dl>

<p>Thus in order to have a spontaneous reaction (<span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle \Delta G_{cell}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle \Delta G_{cell}^{\ominus }}</annotation>

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</math></span><img src="./2b000458eb6269178d0ef5e50f33d5cc4c89d385.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:6.453ex; height:3.176ex;" alt="{\displaystyle \Delta G_{cell}^{\ominus }}" loading="lazy"></span> < 0), <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle E_{cell}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle E_{cell}^{\ominus }}</annotation>

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</math></span><img src="./53586c734844c240fa1837ba007fedf940cf37a3.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:4.406ex; height:3.176ex;" alt="{\displaystyle E_{cell}^{\ominus }}" loading="lazy"></span>must be positive, where:

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<annotation encoding="application/x-tex">{\displaystyle E_{cell}^{\ominus }=E_{cathode}^{\ominus }-E_{anode}^{\ominus }}</annotation>

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</math></span><img src="./2b7b2793f4e32847a9362acdce04c3f78d5f2957.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:24.064ex; height:3.176ex;" alt="{\displaystyle E_{cell}^{\ominus }=E_{cathode}^{\ominus }-E_{anode}^{\ominus }}" loading="lazy"></span></dd></dl>

<p>where <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle E_{cathode}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle E_{cathode}^{\ominus }}</annotation>

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</math></span><img src="./20c5fb08e7b3a6985a949af7a90b5877d3630a39.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:7.493ex; height:3.176ex;" alt="{\displaystyle E_{cathode}^{\ominus }}" loading="lazy"></span> is the standard potential at the cathode (called as standard cathodic potential or standard reduction potential and <span class="mwe-math-element mwe-math-element-inline"><span class="mwe-math-mathml-inline mwe-math-mathml-a11y" style="display: none;"><math xmlns="http://www.w3.org/1998/Math/MathML" alttext="{\displaystyle E_{anode}^{\ominus }}">

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<annotation encoding="application/x-tex">{\displaystyle E_{anode}^{\ominus }}</annotation>

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</math></span><img src="./b5485b261a0bda0a9e8a15f5983268416efb527b.svg" class="mwe-math-fallback-image-inline mw-invert skin-invert" aria-hidden="true" style="vertical-align: -1.005ex; width:6.227ex; height:3.176ex;" alt="{\displaystyle E_{anode}^{\ominus }}" loading="lazy"></span> is the standard potential at the anode (called as standard anodic potential or standard oxidation potential) as given in the <a href="Standard_electrode_potential_(data_page)" title="Standard electrode potential (data page)">table of standard electrode potential</a>.

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<ul><li><a href="Nernst_equation" title="Nernst equation">Nernst equation</a></li>

<li><a href="Pourbaix_diagram" title="Pourbaix diagram">Pourbaix diagram</a></li>

<li><a href="Solvated_electron" title="Solvated electron">Solvated electron</a></li>

<li><a href="Standard_electrode_potential_(data_page)" title="Standard electrode potential (data page)">Standard electrode potential (data page)</a></li>

<li><a href="Standard_hydrogen_electrode" title="Standard hydrogen electrode">Standard hydrogen electrode</a> (SHE)</li>

<li><a href="Table_of_standard_reduction_potentials_for_half-reactions_important_in_biochemistry" title="Table of standard reduction potentials for half-reactions important in biochemistry">Biochemically relevant redox potentials (data page)</a></li></ul>

<div class="mw-heading mw-heading2"><h2 id="References">References</h2></div>

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<li id="cite_note-1"><span class="mw-cite-backlink"><b><a href="#cite_ref-1">^</a></b></span> <span class="reference-text"><a href="International_Union_of_Pure_and_Applied_Chemistry" title="International Union of Pure and Applied Chemistry">IUPAC</a>, <i><a href="IUPAC_books" class="mw-redirect" title="IUPAC books">Compendium of Chemical Terminology</a></i>, 5th ed. (the "Gold Book") (2025). Online version: (2006–) "<a rel="nofollow" class="external text" href="https://goldbook.iupac.org/terms/view/S05912.html">Standard electrode potential, E⚬</a>". <style data-mw-deduplicate="TemplateStyles:r1238218222">

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</style><a href="Doi_(identifier)" class="mw-redirect" title="Doi (identifier)">doi</a>:<a rel="nofollow" class="external text" href="https://doi.org/10.1351%2Fgoldbook.S05912">10.1351/goldbook.S05912</a></span>

</li>

<li id="cite_note-2"><span class="mw-cite-backlink"><b><a href="#cite_ref-2">^</a></b></span> <span class="reference-text"><a rel="nofollow" class="external text" href="http://goldbook.iupac.org/E01956.html">IUPAC definition of the electrode potential</a></span>

</li>

</ol></div></div>

<div class="mw-heading mw-heading2"><h2 id="Further_reading">Further reading</h2></div>

<ul><li>Zumdahl, Steven S., Zumdahl, Susan A (2000) <i>Chemistry</i> (5th ed.), Houghton Mifflin Company. <a href="ISBN_(identifier)" class="mw-redirect" title="ISBN (identifier)">ISBN</a> <bdi>0-395-98583-8</bdi></li>

<li>Atkins, Peter, Jones, Loretta (2005) <i>Chemical Principles</i> (3rd ed.), W.H. Freeman and Company. <a href="ISBN_(identifier)" class="mw-redirect" title="ISBN (identifier)">ISBN</a> <bdi>0-7167-5701-X</bdi></li>

<li>Zu, Y, Couture, MM, Kolling, DR, Crofts, AR, Eltis, LD, Fee, JA, Hirst, J (2003) <i>Biochemistry</i>, 42, 12400-12408</li></ul>

<div class="mw-heading mw-heading2"><h2 id="External_links">External links</h2></div>

<ul><li><a rel="nofollow" class="external text" href="http://hyperphysics.phy-astr.gsu.edu/hbase/Tables/electpot.html">Standard Electrode Potentials table</a></li>

<li><a rel="nofollow" class="external text" href="http://www.chemguide.co.uk/physical/redoxeqia/introduction.html">Redox Equilibria</a></li>

<li><a rel="nofollow" class="external text" href="http://www.science.uwaterloo.ca/~cchieh/cact/c123/battery.html">Chemistry of Batteries</a></li>

<li><a rel="nofollow" class="external text" href="http://hyperphysics.phy-astr.gsu.edu/HBASE/Chemical/electrochem.html#c1">Electrochemical Cells</a></li>

<li><a rel="nofollow" class="external text" href="http://www.tandfonline.com/doi/abs/10.1080/14786440908564891">STEP in Non-aqueous solvent</a></li></ul></div><div><div class="zim-footer">

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