Table of standard reduction potentials (electrochemical series)

Standard reduction potentials, E°, for common half-reactions at 25 °C, from the strongest oxidizing agent at the top to the strongest reducing agent at the bottom. Pick two half-reactions to get the voltage of a cell.

Cell voltage calculator

E°cell = E°cathode − E°anode = +0.34 − (−0.76) +1.10 V Positive, so the reaction runs on its own: a galvanic cell.

Standard reduction potentials at 25 °C

Strongest oxidizing agents at the top, strongest reducing agents at the bottom. Tap a symbol to open the element.

Half-reaction (reduction)E° (V)E° on a scale from −3 V to +3 V
F F₂ + 2e⁻ → 2F⁻ +2.87
O H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O +1.78
Mn MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O +1.51
Au Au³⁺ + 3e⁻ → Au +1.50
Cl Cl₂ + 2e⁻ → 2Cl⁻ +1.36
Cr Cr₂O₇²⁻ + 14H⁺ + 6e⁻ → 2Cr³⁺ + 7H₂O +1.33
O O₂ + 4H⁺ + 4e⁻ → 2H₂O +1.23
Mn MnO₂ + 4H⁺ + 2e⁻ → Mn²⁺ + 2H₂O +1.22
Pt Pt²⁺ + 2e⁻ → Pt +1.18
Br Br₂ + 2e⁻ → 2Br⁻ +1.07
N NO₃⁻ + 4H⁺ + 3e⁻ → NO + 2H₂O +0.96
Hg Hg²⁺ + 2e⁻ → Hg +0.85
Ag Ag⁺ + e⁻ → Ag +0.80
Fe Fe³⁺ + e⁻ → Fe²⁺ +0.77
O O₂ + 2H⁺ + 2e⁻ → H₂O₂ +0.70
I I₂ + 2e⁻ → 2I⁻ +0.54
Cu Cu⁺ + e⁻ → Cu +0.52
O O₂ + 2H₂O + 4e⁻ → 4OH⁻ +0.40
Cu Cu²⁺ + 2e⁻ → Cu +0.34
Ag AgCl + e⁻ → Ag + Cl⁻ +0.22
Cu Cu²⁺ + e⁻ → Cu⁺ +0.15
Sn Sn⁴⁺ + 2e⁻ → Sn²⁺ +0.15
H 2H⁺ + 2e⁻ → H₂ 0.00
Fe Fe³⁺ + 3e⁻ → Fe −0.04
Pb Pb²⁺ + 2e⁻ → Pb −0.13
Sn Sn²⁺ + 2e⁻ → Sn −0.14
Ni Ni²⁺ + 2e⁻ → Ni −0.26
Co Co²⁺ + 2e⁻ → Co −0.28
Cd Cd²⁺ + 2e⁻ → Cd −0.40
Fe Fe²⁺ + 2e⁻ → Fe −0.45
Cr Cr³⁺ + 3e⁻ → Cr −0.74
Zn Zn²⁺ + 2e⁻ → Zn −0.76
H 2H₂O + 2e⁻ → H₂ + 2OH⁻ −0.83
Mn Mn²⁺ + 2e⁻ → Mn −1.19
Al Al³⁺ + 3e⁻ → Al −1.66
Mg Mg²⁺ + 2e⁻ → Mg −2.37
Na Na⁺ + e⁻ → Na −2.71
Ca Ca²⁺ + 2e⁻ → Ca −2.87
Ba Ba²⁺ + 2e⁻ → Ba −2.91
K K⁺ + e⁻ → K −2.93
Li Li⁺ + e⁻ → Li −3.04

How to read the table

Each line is a reduction half-reaction, in which the species on the left gains electrons. The more positive E° is, the more strongly that species pulls electrons in, so fluorine, at the top, is the strongest oxidizing agent listed. The more negative E° is, the more readily the species on the right gives electrons away, so lithium metal, at the bottom, is the strongest reducing agent.

Every value is measured against the standard hydrogen electrode, which is defined as exactly 0 V, under standard conditions: 25 °C, gases at 1 bar and dissolved species at 1 mol/L.

Working out a cell voltage

In a galvanic cell, reduction happens at the cathode and oxidation at the anode. The cell potential is E°cell = E°cathode − E°anode, using both values exactly as the table gives them. Do not flip the sign of the anode value yourself, and do not multiply E° by the number of electrons: it does not depend on how many times a half-reaction is written.

Example: the Daniell cell pairs copper with zinc. E°cell = +0.34 V − (−0.76 V) = +1.10 V. A positive E°cell means the reaction runs on its own as written.

Predicting which reactions happen

The species on the left of any line can oxidize the species on the right of any line below it. Copper(II) ions (+0.34 V) take electrons from zinc metal (−0.76 V), so a zinc strip dipped in copper sulfate solution becomes coated with copper. Silver ions (+0.80 V) are reduced by copper metal, which is why copper wire in silver nitrate solution grows silver crystals.

Metals with negative E° values, below hydrogen, can in principle release hydrogen gas from acids, while copper, silver and gold, above it, cannot. This matches the order of the reactivity series of metals.

Why values differ between books

Sources round E° values differently and some use older measurements, so iron(II) may appear as −0.44 V or −0.45 V. The values here come from standard reference data and are shown to two decimal places; use the numbers your course gives when they differ.

FAQ

What is a standard reduction potential?

It is the voltage of a reduction half-reaction measured against the standard hydrogen electrode at 25 °C, 1 bar and 1 mol/L. It shows how strongly a species attracts electrons: the higher the value, the more easily it is reduced.

What is the strongest oxidizing agent?

Fluorine, F₂, with E° = +2.87 V, is the strongest common oxidizing agent. Lithium metal, at −3.04 V, is the strongest common reducing agent.

How do you calculate E°cell?

Subtract the anode potential from the cathode potential: E°cell = E°cathode − E°anode. For a zinc–copper cell, +0.34 − (−0.76) = +1.10 V.

Why is the hydrogen electrode 0 V?

Only differences in potential can be measured, so chemists chose the standard hydrogen electrode, 2H⁺ + 2e⁻ → H₂, as the reference and defined its potential as exactly 0 V.