Standard Reduction Potentials (E°) Table
| Half-Reaction (Reduction) | E° (V) | Redox Couple | Category |
|---|---|---|---|
| Li⁺(aq) + e⁻ → Li(s) | -3.04 | Li⁺/Li | Alkali metal |
| K⁺(aq) + e⁻ → K(s) | -2.924 | K⁺/K | Alkali metal |
| Ca²⁺(aq) + 2e⁻ → Ca(s) | -2.868 | Ca²⁺/Ca | Alkaline earth |
| Na⁺(aq) + e⁻ → Na(s) | -2.714 | Na⁺/Na | Alkali metal |
| Mg²⁺(aq) + 2e⁻ → Mg(s) | -2.372 | Mg²⁺/Mg | Alkaline earth |
| Al³⁺(aq) + 3e⁻ → Al(s) | -1.662 | Al³⁺/Al | Post-transition metal |
| Mn²⁺(aq) + 2e⁻ → Mn(s) | -1.185 | Mn²⁺/Mn | Transition metal |
| Zn²⁺(aq) + 2e⁻ → Zn(s) | -0.762 | Zn²⁺/Zn | Transition metal |
| Cr³⁺(aq) + 3e⁻ → Cr(s) | -0.744 | Cr³⁺/Cr | Transition metal |
| Fe²⁺(aq) + 2e⁻ → Fe(s) | -0.447 | Fe²⁺/Fe | Transition metal |
| Cr³⁺(aq) + e⁻ → Cr²⁺(aq) | -0.407 | Cr³⁺/Cr²⁺ | Transition metal |
| Co²⁺(aq) + 2e⁻ → Co(s) | -0.28 | Co²⁺/Co | Transition metal |
| Ni²⁺(aq) + 2e⁻ → Ni(s) | -0.257 | Ni²⁺/Ni | Transition metal |
| Sn²⁺(aq) + 2e⁻ → Sn(s) | -0.138 | Sn²⁺/Sn | Post-transition metal |
| Pb²⁺(aq) + 2e⁻ → Pb(s) | -0.126 | Pb²⁺/Pb | Post-transition metal |
| 2H⁺(aq) + 2e⁻ → H₂(g) | 0 | H⁺/H₂ | Reference (SHE) |
| Sn⁴⁺(aq) + 2e⁻ → Sn²⁺(aq) | 0.151 | Sn⁴⁺/Sn²⁺ | Post-transition metal |
| Cu²⁺(aq) + e⁻ → Cu⁺(aq) | 0.153 | Cu²⁺/Cu⁺ | Transition metal |
| Cu²⁺(aq) + 2e⁻ → Cu(s) | 0.342 | Cu²⁺/Cu | Transition metal |
| I₂(s) + 2e⁻ → 2I⁻(aq) | 0.536 | I₂/I⁻ | Halogen |
| Fe³⁺(aq) + e⁻ → Fe²⁺(aq) | 0.771 | Fe³⁺/Fe²⁺ | Transition metal |
| Ag⁺(aq) + e⁻ → Ag(s) | 0.8 | Ag⁺/Ag | Transition metal |
| Hg²⁺(aq) + 2e⁻ → Hg(l) | 0.851 | Hg²⁺/Hg | Transition metal |
| Br₂(l) + 2e⁻ → 2Br⁻(aq) | 1.066 | Br₂/Br⁻ | Halogen |
| O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) | 1.229 | O₂/H₂O | Nonmetal |
| Cr₂O₇²⁻(aq) + 14H⁺(aq) + 6e⁻ → 2Cr³⁺(aq) + 7H₂O(l) | 1.232 | Cr₂O₇²⁻/Cr³⁺ | Transition metal |
| Cl₂(g) + 2e⁻ → 2Cl⁻(aq) | 1.358 | Cl₂/Cl⁻ | Halogen |
| Au³⁺(aq) + 3e⁻ → Au(s) | 1.498 | Au³⁺/Au | Transition metal |
| MnO₄⁻(aq) + 8H⁺(aq) + 5e⁻ → Mn²⁺(aq) + 4H₂O(l) | 1.507 | MnO₄⁻/Mn²⁺ | Transition metal |
| PbO₂(s) + 4H⁺(aq) + SO₄²⁻(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l) | 1.685 | PbO₂/PbSO₄ | Lead-acid battery |
| H₂O₂(aq) + 2H⁺(aq) + 2e⁻ → 2H₂O(l) | 1.776 | H₂O₂/H₂O | Nonmetal |
| Co³⁺(aq) + e⁻ → Co²⁺(aq) | 1.92 | Co³⁺/Co²⁺ | Transition metal |
| F₂(g) + 2e⁻ → 2F⁻(aq) | 2.866 | F₂/F⁻ | Halogen |
All potentials are tabulated at 298.15 K, 1 atm gas pressure, and 1 M for aqueous species, referenced to the SHE. Caveat: real systems rarely sit at standard concentrations — use the Nernst equation (E = E° − (RT/nF) ln Q) to correct for actual ion concentrations. E° does not depend on the stoichiometric coefficients of the half-reaction; multiplying a half-reaction by 2 doesn't double E°. Some half-reactions are pH-dependent (those with H⁺ in the equation), and the potentials listed assume [H⁺] = 1 M. Sources: CRC Handbook of Chemistry and Physics and the NIST Standard Reference Database.
Frequently Asked Questions
How do you calculate standard cell potential from the table?
Identify which half-reaction runs as reduction (cathode) and which runs as oxidation (anode), look up the reduction potential for each as written in the table, then compute E°cell = E°cathode − E°anode. Don't flip the sign of the anode value before subtracting — the formula already accounts for the reversal. For the Daniell cell with Zn|Zn²⁺ at the anode and Cu²⁺|Cu at the cathode: E°cell = (+0.342) − (−0.762) = +1.104 V. Positive means the reaction is spontaneous under standard conditions.
Why is the standard hydrogen electrode set to exactly 0 V?
Absolute electrode potentials can't be measured directly — only the difference between two electrodes is observable. By convention IUPAC fixes the SHE at exactly 0.000 V to anchor the entire scale. The half-reaction is 2H⁺(aq, 1 M) + 2e⁻ → H₂(g, 1 atm) at 25 °C, with a platinized platinum electrode. Every other reduction potential in the table is the cell voltage measured against this reference. Choosing a different reference (like the Ag/AgCl electrode at +0.197 V) would shift every value by the same constant — the differences, which determine cell behavior, would be unchanged.
What does a more negative reduction potential mean?
A more negative E° means the oxidized form on the left is harder to reduce — it doesn't want the electrons. The flip side is that the reduced form on the right is a strong electron donor, i.e., a strong reducing agent. Li⁺ + e⁻ → Li sits at −3.04 V, so Li⁺ resists reduction but Li metal eagerly gives up its electron, making it one of the strongest reducing agents. The opposite extreme is F₂ + 2e⁻ → 2F⁻ at +2.87 V: F₂ is the strongest common oxidizing agent because the reduction is wildly favorable.