Every electrochemical couple has a characteristic tendency to gain or give up electrons. The standard electrode potential, E°, puts a number on that tendency so that different half-reactions can be ranked on one scale. It is the starting point for estimating the voltage of any battery chemistry, and for understanding why some materials make good anodes while others make good cathodes.
What the number means
A single electrode potential cannot be measured on its own: a voltmeter always sees the difference between two electrodes. IUPAC therefore defines the standard electrode potential as the standard emf of a cell in which the electrode of interest is paired with a hydrogen electrode, where hydrogen at standard pressure is oxidised to solvated protons [3]. The standard hydrogen electrode (SHE) is assigned 0 V, and every other half-reaction is quoted relative to it, written as a reduction, at 25 °C with all species at unit activity.
A large negative value means the metal strongly prefers to be oxidised (it is a strong reducing agent), which is what you want at the negative electrode. A large positive value means the oxidised species strongly prefers to take electrons, which is what you want at the positive electrode.
Estimating a cell voltage
Under standard conditions the ideal open circuit voltage of a cell is the difference between the positive and negative half-reaction potentials:
The lead acid cell is a useful worked example because both electrodes appear in the table below. The positive plate couple (PbO₂/PbSO₄) sits at +1.6913 V and the negative plate couple (PbSO₄/Pb) sits at −0.3588 V [1], giving a standard cell voltage of about 2.05 V. That matches the familiar ~2 V per cell of a lead acid battery.
Real cells deviate from E° because concentrations, temperature and state of charge move the potentials; the Nernst Equation describes that correction. The measured Open Circuit Voltage is the practical quantity a battery engineer works with.
Table of battery-relevant standard reduction potentials
Values are at 25 °C versus SHE. The primary column is taken from the Electrochemical Series in the CRC Handbook of Chemistry and Physics [1]. The second column is an independent compilation [2], included as a cross-check; rows marked * differ between the two by 10 mV or more, and a dash means that couple is not listed in [2]. Small differences are normal: the compilations draw on different underlying datasets, and [1] notes that values for less common reactions are subject to revision. The fourth column adds 3.0401 V to the CRC value to re-express it against the Li⁺/Li couple, which is the scale used throughout lithium-ion and most non-aqueous battery literature.
| Half-reaction (reduction) | E° vs SHE (V) [1] | E° vs SHE (V) [2] | E° vs Li⁺/Li (V) | Battery relevance |
|---|---|---|---|---|
| Li⁺ + e⁻ ⇌ Li | -3.0401 | -3.040 | 0.00 | Lithium-ion reference electrode; lithium metal anodes |
| K⁺ + e⁻ ⇌ K | -2.931 | -2.93 | 0.11 | Potassium-ion |
| Ca²⁺ + 2e⁻ ⇌ Ca | -2.868 | -2.84 * | 0.17 | Calcium-ion |
| Na⁺ + e⁻ ⇌ Na | -2.71 | -2.713 | 0.33 | Sodium-ion |
| Mg²⁺ + 2e⁻ ⇌ Mg | -2.372 | -2.356 * | 0.67 | Magnesium-ion (research stage) |
| Al³⁺ + 3e⁻ ⇌ Al | -1.662 | -1.676 * | 1.38 | Aluminium-ion |
| ZnO + H₂O + 2e⁻ ⇌ Zn + 2OH⁻ | -1.260 | – | 1.78 | Zinc anode in alkaline cells (Nickel-Zinc, Zinc-Air, Silver-Zinc) |
| 2H₂O + 2e⁻ ⇌ H₂ + 2OH⁻ | -0.8277 | -0.828 | 2.21 | Hydrogen evolution in alkaline electrolyte |
| Cd(OH)₂ + 2e⁻ ⇌ Cd + 2OH⁻ † | -0.809 | – | 2.23 | Nickel-Cadmium negative electrode |
| Zn²⁺ + 2e⁻ ⇌ Zn | -0.7618 | -0.7618 | 2.28 | Rechargeable aqueous zinc-ion |
| Fe²⁺ + 2e⁻ ⇌ Fe | -0.447 | -0.44 | 2.59 | Iron-Air |
| PbSO₄ + 2e⁻ ⇌ Pb + SO₄²⁻ | -0.3588 | -0.356 | 2.68 | Lead acid negative plate |
| Ni²⁺ + 2e⁻ ⇌ Ni | -0.257 | -0.257 | 2.78 | Nickel metal (current collectors, plating) |
| Pb²⁺ + 2e⁻ ⇌ Pb | -0.1262 | -0.126 | 2.91 | Lead metal |
| 2H⁺ + 2e⁻ ⇌ H₂ | 0.0000 | 0.00 | 3.04 | Standard hydrogen electrode (reference point) |
| Cu²⁺ + 2e⁻ ⇌ Cu | +0.3419 | +0.3419 | 3.38 | Copper current collector dissolution risk |
| Ag₂O + H₂O + 2e⁻ ⇌ 2Ag + 2OH⁻ | +0.342 | – | 3.38 | Silver-Zinc positive electrode |
| O₂ + 2H₂O + 4e⁻ ⇌ 4OH⁻ | +0.401 | +0.401 | 3.44 | Air cathode in alkaline metal-air cells |
| NiO₂ + 2H₂O + 2e⁻ ⇌ Ni(OH)₂ + 2OH⁻ | +0.490 | – | 3.53 | Approximates the nickel positive electrode (NiOOH/Ni(OH)₂) in alkaline nickel chemistries |
| Fe³⁺ + e⁻ ⇌ Fe²⁺ | +0.771 | +0.771 | 3.81 | Iron redox couple (iron-based flow batteries) |
| Ag⁺ + e⁻ ⇌ Ag | +0.7996 | +0.7996 | 3.84 | Silver metal |
| Br₂(aq) + 2e⁻ ⇌ 2Br⁻ | +1.0873 | +1.087 | 4.13 | Hydrogen-Bromine flow battery |
| MnO₂ + 4H⁺ + 2e⁻ ⇌ Mn²⁺ + 2H₂O | +1.224 | +1.23 | 4.26 | Manganese dioxide cathodes |
| O₂ + 4H⁺ + 4e⁻ ⇌ 2H₂O | +1.229 | +1.229 | 4.27 | Oxygen evolution in acidic electrolyte |
| PbO₂ + SO₄²⁻ + 4H⁺ + 2e⁻ ⇌ PbSO₄ + 2H₂O | +1.6913 | +1.690 | 4.73 | Lead acid positive plate |
| F₂ + 2e⁻ ⇌ 2F⁻ | +2.866 | +2.87 | 5.91 | Upper end of the series; relevant to Fluoride-ion concepts |
Reading the table as a battery designer
- Why lithium: Li⁺/Li sits at the very bottom of the series, so pairing it with almost any positive electrode gives the largest available voltage. Combined with lithium’s low mass, this is the root of lithium-based chemistries’ energy density advantage.
- The water limit: the gap between hydrogen evolution (0 V) and oxygen evolution (+1.229 V) in acid is only about 1.23 V. Aqueous chemistries that exceed this rely on kinetic overpotentials to avoid splitting water, which is why lead acid can run at ~2 V per cell but high-voltage chemistries need non-aqueous electrolytes. See Electrochemical Potential Window of Battery Electrolytes.
- Beyond-lithium options: sodium, potassium and calcium are all within about 0.35 V of lithium, so their cells can in principle reach similar voltages, although mass, ion size and electrode chemistry change the picture considerably.
- Alkaline chemistries: pairing the nickel and cadmium couples gives roughly 0.49 − (−0.81) ≈ 1.3 V, in line with the ~1.2 V nominal of NiCd cells, and swapping cadmium for zinc (ZnO/Zn) lifts that towards 1.75 V, which is the idea behind nickel-zinc.
- Copper current collectors: copper’s positive potential is one reason it is only used on the negative electrode side; if the anode is over-discharged to a high potential the copper foil can begin to dissolve.
Limits of standard potentials
Standard potentials describe metal/ion and simple redox couples in aqueous solution. Intercalation electrodes, such as graphite anodes and layered oxide or phosphate cathodes, do not have a single fixed E°: their potential changes continuously with the amount of lithium stored, which is why an OCV curve is needed to describe them. For those materials the Li⁺/Li scale is used directly and the values come from measurement rather than from a table like this one. See Anode and Cathode for the materials themselves.
References
- Vanýsek, P., Electrochemical Series, in CRC Handbook of Chemistry and Physics, CRC Press/Taylor & Francis (Section 8). The tables themselves draw on Milazzo, Caroli and Sharma (1978), Bard, Parsons and Jordan (1985) and Bratsch, J. Phys. Chem. Ref. Data 18 (1989).
- P2: Standard Reduction Potentials by Value, Chemistry LibreTexts (D. Harvey)
- Standard electrode potential (S05912), IUPAC Compendium of Chemical Terminology (Gold Book)
See also: Galvanic vs Electrolytic, Battery Chemistry Definitions & Glossary