What is a standard reduction potential?

It is the voltage a half-reaction produces, written as a reduction (gain of electrons), measured against the Standard Hydrogen Electrode under standard conditions: 25°C, 1 M solutions, and 1 atm gas. The SHE itself is defined as exactly 0 V, so all other potentials are relative to it.

How do I find a cell's EMF?

Use E°cell = E°cathode − E°anode, where both values are the standard reduction potentials from the table. If the result is positive the cell is galvanic and the reaction proceeds spontaneously; if negative, energy must be supplied as in electrolysis.

What does a high positive E° mean?

A high positive E° marks a strong oxidising agent — the species readily gains electrons and is reduced. Fluorine (+2.87 V) is the strongest common oxidiser. Strongly negative E° values, like lithium (−3.04 V), mark strong reducing agents that readily give up electrons.

Why is hydrogen set to zero volts?

Absolute electrode potentials cannot be measured directly, only differences between electrodes. Chemists adopted the Standard Hydrogen Electrode as the universal reference and assigned it exactly 0 V, so every tabulated potential is the voltage of that half-cell paired against the SHE.

Does the series predict displacement reactions?

Yes. A metal with a more negative reduction potential will displace a metal ion with a less negative potential from solution. Zinc (−0.76 V) displaces copper (+0.34 V) from copper sulfate because the combined cell EMF is positive, the basis of the reactivity series.

What is the Electrochemical Series Reference?

Reference table of standard reduction potentials (E° in volts vs SHE) for 35+ common half-reactions, plus a cell-EMF calculator that tells you whether a galvanic cell is spontaneous. Runs in your browser. It runs free in your browser on Gera Tools, with nothing uploaded.

Electrochemical Series Reference

Name: Electrochemical Series Reference
Creator: Gera Tools
License: https://creativecommons.org/licenses/by/4.0/

The electrochemical series ranks half-reactions by their standard reduction potential, the voltage each one produces against a common reference. This reference lists E° values for more than thirty-five common half-reactions and includes a cell-EMF calculator so you can check whether a galvanic cell is spontaneous.

How it works

Every half-reaction is written as a reduction — electrons on the left — with its standard electrode potential E° measured in volts against the Standard Hydrogen Electrode (SHE), which is defined as exactly 0 V at 25°C, 1 M, and 1 atm. To find the voltage of a complete cell, combine two half-reactions:

E°cell = E°cathode − E°anode

The cathode is where reduction occurs and the anode where oxidation occurs. If E°cell is positive the reaction is spontaneous and the cell is galvanic; if negative, it needs external energy, as in electrolysis.

Reading the series

Species near the top have large positive potentials and are strong oxidising agents — they pull electrons toward themselves and are easily reduced. Fluorine sits at the top at +2.87 V. Species near the bottom have strongly negative potentials and are strong reducing agents, readily giving up electrons; lithium anchors the bottom at −3.04 V. The position of a metal also predicts displacement reactions and underpins the familiar reactivity series.

Selected standard reduction potentials

A quick reference for commonly tested and practically important half-reactions:

Half-reaction (reduction)	E° (V vs SHE)	Significance
F₂ + 2e⁻ → 2F⁻	+2.87	Strongest common oxidising agent
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O	+1.51	Permanganate oxidation
Cl₂ + 2e⁻ → 2Cl⁻	+1.36	Chlorine as oxidant / bleach
O₂ + 4H⁺ + 4e⁻ → 2H₂O	+1.23	Oxygen reduction in fuel cells
Ag⁺ + e⁻ → Ag	+0.80	Silver plating
Cu²⁺ + 2e⁻ → Cu	+0.34	Copper deposition
2H⁺ + 2e⁻ → H₂	0.00	Standard Hydrogen Electrode reference
Pb²⁺ + 2e⁻ → Pb	−0.13	Lead-acid battery negative plate
Ni²⁺ + 2e⁻ → Ni	−0.25	Nickel plating / galvanic corrosion
Fe²⁺ + 2e⁻ → Fe	−0.44	Iron rusting (anodic reaction)
Zn²⁺ + 2e⁻ → Zn	−0.76	Zinc galvanizing protects steel
Al³⁺ + 3e⁻ → Al	−1.66	Aluminium: light and anodically active
Li⁺ + e⁻ → Li	−3.04	Strongest common reducing agent

Worked example: Daniell cell

A Daniell cell pairs the copper cathode (Cu²⁺ + 2e⁻ → Cu, +0.34 V) with a zinc anode (Zn²⁺ + 2e⁻ → Zn, −0.76 V). Its EMF is 0.34 − (−0.76) = +1.10 V, which is positive, so the cell is spontaneous — exactly what a real Daniell cell delivers.

The overall cell reaction writes the anode as oxidation (reverse the reduction): Zn → Zn²⁺ + 2e⁻, then adds it to the cathode reduction: Zn + Cu²⁺ → Zn²⁺ + Cu. Zinc displaces copper from solution because zinc’s reduction potential is more negative — it is the better reducing agent.

Practical applications of the series

Corrosion prediction. When two dissimilar metals are in electrical contact in an electrolyte, the one with the more negative E° acts as the anode and corrodes preferentially. Zinc corrodes instead of steel in galvanized fencing for exactly this reason.
Battery design. A high cell voltage requires the largest possible gap between cathode and anode potentials. Lithium batteries achieve very high voltages because lithium sits at −3.04 V while the cathode materials are near +1 V.
Electrolysis. If E°cell for a desired reaction is negative, you must supply at least that voltage to drive it. Electrolytic production of aluminium requires voltages around 4 V because aluminium’s reduction potential is so negative.

All potentials here are standard values at 25°C and 1 M. In practice the Nernst equation adjusts cell voltage for non-standard concentrations and temperature.