# Standard electrode potential The **standard electrode potential**, denoted \( E^{\circ} \), is the standard electromotive force that a cell would have when, at its left-hand electrode, molecular hydrogen at standard pressure undergoes oxidation to give solvated protons.[1](https://goldbook.iupac.org/terms/view/S05912) It is the internationally agreed measure of the tendency of a half-reaction to occur as a reduction, expressed on a scale whose zero point is the standard hydrogen electrode[2](https://goldbook.iupac.org/terms/view/S05917). IUPAC treats "standard potential of an electrode reaction" as synonymous with the term[3](https://goldbook.iupac.org/terms/view/S05919). | Key fact | Detail | |---|---| | Definition | Standard emf of a cell with molecular hydrogen under standard pressure oxidized to solvated protons at the left-hand electrode[1](https://goldbook.iupac.org/terms/view/S05912) | | Zero point | The standard hydrogen electrode, for which \( E^{\circ}(\ce{H+}/\ce{H2}) \) is zero at all temperatures[2](https://goldbook.iupac.org/terms/view/S05917) | | Absolute anchor | Recommended absolute potential of the hydrogen electrode in water: \( (4.44 \pm 0.02) \ \mathrm{V} \) at 298.15 K[2](https://goldbook.iupac.org/terms/view/S05917) | | Sign convention | Electrode potentials are reduction potentials; the electrode in question sits on the right of the cell diagram with the SHE on the left[4](https://goldbook.iupac.org/terms/view/E01956) | | Thermodynamic link | \( \Delta G^{\circ} = -nFE^{\circ} \), with \( F = 96{,}485 \ \mathrm{C/mol\ e^-} \)[5](https://openstax.org/books/chemistry-2e/pages/17-4-potential-free-energy-and-equilibrium) | | Concentration dependence | Nernst equation \( E_{\mathrm{eq}} = E^{\circ} - (RT/zF)\sum_i \nu_i \ln(a_i) \)[6](https://goldbook.iupac.org/terms/view/09068) | | Tabulated coverage | Standard potentials and temperature coefficients in water at 298.15 K given for nearly 1700 half-reactions at pH 0.000 and pH 13.996[7](https://srd.nist.gov/JPCRD/jpcrd355.pdf) | ## Definition and the reference scale A single electrode's potential cannot be measured in isolation; the absolute value of an electrode potential cannot be measured, so \( E \) is always reported relative to the potential of a reference electrode, typically the standard hydrogen electrode (SHE)[8](https://iupac.org/wp-content/uploads/2019/07/PAC-REC-18-01-09R2_PR190703MC.pdf). According to IUPAC, the electrode potential \( E \) is the electromotive force of a [galvanic cell](https://www.edgepedia.org/galvanic-cell) whose left-hand electrode is a standard hydrogen electrode and whose right-hand electrode is the one being considered[4](https://goldbook.iupac.org/terms/view/E01956). The SHE supplies the arbitrary zero. For solutions in protic solvents it is the universal reference electrode for which, under standard conditions, the standard electrode potential of the \( \ce{H+}/\ce{H2} \) couple is zero at all temperatures[2](https://goldbook.iupac.org/terms/view/S05917). In the common textbook construction, the SHE consists of hydrogen gas at 1 atm bubbled through a 1 M HCl solution with an inert platinum electrode, and the reference half-reaction \( 2\ce{H+} + 2\ce{e-} \rightleftharpoons \ce{H2} \) is assigned \( E^{\circ} = 0 \ \mathrm{V} \)[9](https://openstax.org/books/chemistry/pages/17-3-standard-reduction-potentials). **Standard conditions.** The superscript degree sign denotes standard states: 1 bar or 1 atm for gases and 1 M for solutes, often at 298.15 K[9](https://openstax.org/books/chemistry/pages/17-3-standard-reduction-potentials). Potentials measured under unit concentration rather than unit activity are reported as the formal potential \( E^{\circ\prime} \), which can differ from \( E^{\circ} \) owing to real conditions such as pH, ionic strength, and the concentration of complex-forming substances[10](https://goldbook.iupac.org/terms/view/09065). ## Sign convention and reduction potentials By the IUPAC convention, all electrode potentials in tabulations are reduction potentials, meaning the electrode reaction is written in the direction of reduction[11](https://publications.iupac.org/publications/pac/1986/pdf/5807x0955.pdf). A more positive \( E^{\circ} \) therefore means a greater tendency for the oxidized form to gain electrons. When two half-cells are combined, the standard cell potential is the standard reduction potential at the cathode minus that at the anode; for the silver–copper pair, \( E^{\circ}_{\mathrm{cell}} = 0.80 \ \mathrm{V} - 0.34 \ \mathrm{V} = 0.46 \ \mathrm{V} \)[12](https://openstax.org/books/chemistry-atoms-first/pages/16-3-standard-reduction-potentials). Potentials are intensive: when a half-reaction is multiplied by a stoichiometric factor to balance an overall equation, its standard reduction potential is not scaled[12](https://openstax.org/books/chemistry-atoms-first/pages/16-3-standard-reduction-potentials). ## Relation to thermodynamics The standard electrode potential connects directly to the Gibbs energy change of the cell reaction through \( \Delta G^{\circ} = -nFE^{\circ} \), where \( n \) is the number of moles of electrons transferred and \( F \) is the [Faraday constant](https://www.edgepedia.org/faraday-constant), 96,485 C per mole of electrons[5](https://openstax.org/books/chemistry-2e/pages/17-4-potential-free-energy-and-equilibrium). Spontaneous redox reactions under standard conditions show positive cell potentials and negative free energy changes[5](https://openstax.org/books/chemistry-2e/pages/17-4-potential-free-energy-and-equilibrium). Away from standard conditions, the equilibrium potential follows the [Nernst equation](https://www.edgepedia.org/nernst-equation), written as a reduction: \[ E_{\mathrm{eq}} = E^{\circ} - \frac{RT}{zF} \sum_i \nu_i \ln(a_i) \] where \( R \) is the gas constant, \( T \) the thermodynamic temperature, \( F \) the Faraday constant, \( z \) the electron number of the reaction, \( \nu_i \) the stoichiometric coefficients (positive for products, negative for reactants), and \( a_i \) the activities[6](https://goldbook.iupac.org/terms/view/09068). For a redox couple this reduces to \( E_{\mathrm{eq}} = E^{\circ} - (RT/zF)\ln(a_{\mathrm{red}}/a_{\mathrm{ox}}) \)[6](https://goldbook.iupac.org/terms/view/09068). At 25 °C the practical form uses the factor \( 0.0592/z \) with base-10 logarithms[6](https://goldbook.iupac.org/terms/view/09068). ## Tabulated values and examples Because every half-reaction is measured against the same zero, tabulated potentials can be combined to predict the cell potential of any pair of electrodes. A NIST-based compilation gives standard electrode potentials and temperature coefficients in water at 298.15 K for nearly 1700 half-reactions, at pH 0.000 and pH 13.996, drawing primarily on the NBS Tables of Chemical Thermodynamic Properties[7](https://srd.nist.gov/JPCRD/jpcrd355.pdf). Temperature coefficients are defined relative to that of the SHE being zero at all temperatures, and potentials at other temperatures are estimated as \( E_T^{0} = E_{298}^{0} + (T - 298.15) \cdot (dE^{0}/dT)_{298} \)[7](https://srd.nist.gov/JPCRD/jpcrd355.pdf). Selected standard reduction potentials at 298 K illustrate the range of the scale[13](https://openstax.org/books/chemistry/pages/l-standard-electrode-half-cell-potentials): | Half-reaction | \( E^{\circ} \) (V) | |---|---| | \( \ce{F2} + 2\ce{e-} \rightarrow 2\ce{F-} \) | +2.866 | | \( \ce{Cu^2+} + 2\ce{e-} \rightarrow \ce{Cu} \) | +0.34 | | \( 2\ce{H3O+} + 2\ce{e-} \rightarrow \ce{H2} + 2\ce{H2O} \) | 0.00 | | \( \ce{Zn^2+} + 2\ce{e-} \rightarrow \ce{Zn} \) | −0.7618 | | \( \ce{Li+} + \ce{e-} \rightarrow \ce{Li} \) | −3.04 | Fluorine, at the positive end of this selection, is the strongest oxidizing agent listed; lithium metal, at the negative end, is the strongest reducing agent. ## Absolute potential The zero of the conventional scale is a definition, not a measurement. The absolute electrode potential, the potential of a metal measured with respect to a universal reference system without an additional metal/solution interface, can be expressed through a [Born–Haber cycle](https://www.edgepedia.org/born-haber-cycle) involving the atomization and ionization Gibbs energies of \( \ce{H2} \) and the real potential of the proton in the solvent[2](https://goldbook.iupac.org/terms/view/S05917). IUPAC's recommended value for the absolute potential of the standard hydrogen electrode in water is \( (4.44 \pm 0.02) \ \mathrm{V} \) at 298.15 K[2](https://goldbook.iupac.org/terms/view/S05917). Adding this offset converts any tabulated potential versus SHE to the absolute scale. ## What standard potentials do and do not predict A positive standard cell potential predicts that a redox reaction is thermodynamically feasible under standard conditions. This prediction carries two standing qualifications. First, \( E^{\circ} \) values indicate thermodynamic feasibility only and give no information about reaction rate; a feasible reaction may proceed too slowly to observe because of a high activation energy[14](https://simplestudy.com/gb/a-level/ocr/chemistry-a/revision-notes/physical-chemistry-and-transition-elements/redox-and-electrode-potentials). Second, tabulated values apply at standard conditions, and non-standard concentration, pH, pressure, and temperature shift actual electrode potentials through the Nernst relation, so predictions made from \( E^{\circ} \) alone may not hold under real conditions[14](https://simplestudy.com/gb/a-level/ocr/chemistry-a/revision-notes/physical-chemistry-and-transition-elements/redox-and-electrode-potentials). ## References 1. [IUPAC Gold Book: standard electrode potential (S05912)](https://goldbook.iupac.org/terms/view/S05912) 2. [IUPAC Gold Book: standard hydrogen electrode (S05917)](https://goldbook.iupac.org/terms/view/S05917) 3. [IUPAC Gold Book: standard potential of an electrode reaction (S05919)](https://goldbook.iupac.org/terms/view/S05919) 4. [IUPAC Gold Book: electrode potential (E01956)](https://goldbook.iupac.org/terms/view/E01956) 5. [OpenStax Chemistry 2e, 17.4 Potential, Free Energy, and Equilibrium](https://openstax.org/books/chemistry-2e/pages/17-4-potential-free-energy-and-equilibrium) 6. [IUPAC Gold Book: Nernst equation (09068)](https://goldbook.iupac.org/terms/view/09068) 7. [S. G. Bratsch, Standard Electrode Potentials and Temperature Coefficients in Water at 298.15 K, NIST J. Res.](https://srd.nist.gov/JPCRD/jpcrd355.pdf) 8. [IUPAC Recommendations 2019: Terminology of Electrochemical Methods of Analysis](https://iupac.org/wp-content/uploads/2019/07/PAC-REC-18-01-09R2_PR190703MC.pdf) 9. [OpenStax Chemistry, 17.3 Standard Reduction Potentials](https://openstax.org/books/chemistry/pages/17-3-standard-reduction-potentials) 10. [IUPAC Gold Book: formal electrode potential (09065)](https://goldbook.iupac.org/terms/view/09065) 11. [IUPAC PAC 1986: The absolute electrode potential, an explanatory note](https://publications.iupac.org/publications/pac/1986/pdf/5807x0955.pdf) 12. [OpenStax Chemistry: Atoms First, 16.3 Standard Reduction Potentials](https://openstax.org/books/chemistry-atoms-first/pages/16-3-standard-reduction-potentials) 13. [OpenStax Chemistry, Appendix L: Standard Electrode (Half-Cell) Potentials](https://openstax.org/books/chemistry/pages/l-standard-electrode-half-cell-potentials) 14. [SimpleStudy, Redox and Electrode Potentials revision notes (OCR A-Level Chemistry A)](https://simplestudy.com/gb/a-level/ocr/chemistry-a/revision-notes/physical-chemistry-and-transition-elements/redox-and-electrode-potentials) --- *Topic: Encyclopedia › Physical world and mathematics › Chemistry › Chemical principles and methods › Thermodynamics and equilibrium › Chemical equilibrium › Complexation and redox equilibria* *Initially written Sep 17, 2026 · Reviewed: Sep 30, 2026; Oct 15, 2026 · Edited: Sep 30, 2026; Oct 15, 2026 · Last review: Oct 15, 2026* *Copyright 2026 EdgeChat AI, a subsidiary of Biostate AI.* License: Edgepedia Community License 1.0, https://www.edgepedia.org/edgepedia/license